Inorganic Pharmaceutical Chemistry
3rd Class 1st Semester 2012/2013
Reference text:1- Inorganic Medicinal and Pharmaceutical Chemistry
by Block ,Roche Soine and Wilson
2- Textbook of Organic Medicinal and Fharmaceutical Chemistry by
Wilson and Gisvold
----------------------------------------------------------------------Lecture No. 1
Date : 4/10 2012
Dr. Mohammed Hamed
Atomic & Molecular Structural Complexation
atom is composed of a nucleus (about 10-15m in diameter) at the centre which
contains most of the mass of the atom, and orbiting electrons (negatively
charged particles), which have negligible mass.
The nucleus is composed of protons (positively charged fundamental particles),
and neutrons (uncharged fundamental particles).
The radius of an atom is roughly 10,000 times larger than the diameter of the
nucleus, i.e. it is about 10-10m, or 1 angstrom ).
proton
neutron
electron
relative mass
1
1
1/1836
relative charge
+1
0
-1
These figures illustrate the fact that practically all of the mass of the atom is
contained within a very small region
Atoms can be described by two numbers: the atomic number (Z), which is
equal to the number of protons the atom contains, and the mass number (A),
which is equal to the number of protons plus neutrons.
Working out the numbers of protons and neutrons
No of protons = ATOMIC NUMBER of the atom
The atomic number is also given the more descriptive name of proton number.
No of protons + no of neutrons = MASS NUMBER of the atom
The mass number is also called the nucleon number.
This information can be given simply in the form: 19
9
F
How many protons and neutrons has this atom got?
The atomic number counts the number of protons (9); the mass number counts
protons + neutrons (19). If there are 9 protons, there must be 10 neutrons for the
total to add up to 19.
The atomic number is tied to the position of the element in the Periodic
Table and therefore the number of protons defines what sort of element you are
talking about. So if an atom has 8 protons (atomic number = 8), it must be
oxygen. If an atom has 12 protons (atomic number = 12), it must be magnesium.
Isotopes:- The number of neutrons in an atom can vary within small limits. For
example, there are three kinds of carbon atom 12C, 13C and 14C. They all have
the same number of protons, but the number of neutrons varies.
carbon-12
carbon-13
carbon-14
protons
6
6
6
neutrons
6
7
8
mass number
12
13
14
These different atoms of carbon are called isotopes. The fact that they have
varying numbers of neutrons makes no difference whatsoever to the chemical
reactions of the carbon.
Isotopes are atoms which have the same atomic number but different mass
numbers. They have the same number of protons but different numbers of
neutrons
The electrons
Working out the number of electrons
Atoms are electrically neutral, and the positiveness of the protons is balanced by
the negativeness of the electrons. It follows that in a neutral atom:
no of electrons = no of protons
So, if an oxygen atom (atomic number = 8) has 8 protons, it must also have 8
electrons; if a chlorine atom (atomic number = 17) has 17 protons, it must also
have 17 electrons.
How many protons, electrons and neutrons are in an atom of krypton,
carbon, oxygen, neon, silver, gold, etc...?
To find the number of protons, electrons and neutrons in an atom, just
follow these easy steps:
Step 1 - Gather Information
The first thing you will need to do is find some information about your
element. Go to the Periodic Table of Elements .
Use the Table of Elements to find your element's atomic number and
atomic weight. The atomic number is the number located in the upper left
corner and the atomic weight is the number located on the bottom, as in this
example for krypton:
Step 2 - The Number of Protons is...
The atomic number is the number of protons in an atom of an element. In
our example, krypton's atomic number is 36. This tells us that an atom of
krypton has 36 protons in its nucleus.
Step 3 - The Number of Electrons is...
By definition, atoms have no overall electrical charge. That means that
there must be a balance between the positively charged protons and the
negatively charged electrons. Atoms must have equal numbers of protons
and electrons. In our example, an atom of krypton must contain 36
electrons since it contains 36 protons.
Electrons are arranged around atoms in a special way.
An atom can gain or lose electrons, becoming what is known as an ion. An
ion is nothing more than an electrically charged atom. Adding or removing
electrons from an atom does not change which element it is, just its net
charge.
For example, removing an electron from an atom of krypton forms a
krypton ion, which is usually written as Kr+. The plus sign means that this
is a positively charged ion. It is positively charged because a negatively
charged electron was removed from the atom. The 35 remaining electrons
were outnumbered by the 36 positively charged protons, resulting in a
charge of +1.
Step 4 - The Number of Neutrons is...
The atomic weight is basically a measurement of the total number of
particles in an atom's nucleus. In reality, it isn't that clean cut. The atomic
weight is actually a weighted average of all of the naturally
occurring isotopes of an element relative to the mass of carbon-12. You need
to find the mass number. Unfortunately, the mass number isn't listed on
the Table of Elements, to find the mass number, all you need to do is round
the atomic weight to the nearest whole number. In our example,
krypton's mass number is 84 since its atomic weight, 83.80, rounds up to
84.
The mass number is a count of the number of particles in an atom's nucleus.
Remember that the nucleus is made up of protons and neutrons. So, if we
want, we can write:
Mass Number = (Number of Protons) + (Number of Neutrons)
For krypton, this equation becomes:
84 = (Number of Protons) + (Number of Neutrons)
If we only knew how many protons krypton has, we could figure out how
many neutrons it has. Wait a minute... We do know how many protons
krypton has! We did that back in Step 2! The atomic number (36) is the
number of protons in krypton. Putting this into the equation, we get:
84 = 36 + (Number of Neutrons)
What number added to 36 makes 84? Hopefully, you said 48. That is the
number of neutrons in an atom of krypton.
The interesting thing here is that adding or removing neutrons from an atom
does not create a different element. Rather, it creates a heavier or lighter
version of that element. These different versions are called isotopes and
most elements are actually a mixture of different isotopes.
Modern Atomic Theory
Near the end of the 18th century, two laws about chemical reactions emerged
without referring to the notion of an atomic theory. The first was the law of
conservation of mass, formulated by Antoine Lavoisier in 1789, which states
that the total mass in a chemical reaction remains constant (that is, the reactants
have the same mass as the products) The second was the law of definite
proportions. First proven by the French chemist Joseph Louis Proust in
1799, this law states that if a compound is broken down into its constituent
elements, then the masses of the constituents will always have the same
proportions, regardless of the quantity or source of the original substance.
Discovery of subatomic particles
Thomson's illustration of the Crookes tube by which he proved the particle
nature of cathode rays. Cathode rays were emitted from the cathode C,
sharpened to a beam by slits A and B, then passed through the electric field
generated between plates D and E.
When the cathode ray (blue line) passed through the electric field (yellow), it
was deflected.
Atoms were thought to be the smallest possible division of matter until 1897
when J.J. Thomson discovered the electron through his work on cathode rays.
A Crookes tube is a sealed glass container in which two electrodes are separated
by a vacuum. When a voltage is applied across the electrodes, cathode rays are
generated, creating a glowing patch where they strike the glass at the opposite
end of the tube. Through experimentation, Thomson discovered that the rays
could be deflected by an electric field (in addition to magnetic fields, which was
already known). He concluded that these rays, rather than being a form of light,
were composed of very light negatively charged particles he called "corpuscles"
(they would later be renamed electrons by other scientists).
Thomson believed that the corpuscles emerged from the molecules of gas
around the cathode. He thus concluded that atoms were divisible, and that the
corpuscles were their building blocks. To explain the overall neutral charge of
the atom, he proposed that the corpuscles were distributed in a uniform sea of
positive charge; this was the plum pudding model as the electrons were
embedded in the positive charge like plums in a plum pudding (although in
Thomson's model they were not stationary). Thomson's illustration of the
Crookes tube by which he proved the particle nature of cathode rays. Cathode
rays were emitted from the cathode C, sharpened to a beam by slits A and B,
then passed through the electric field generated between plates D and E. When
the cathode ray (blue line) passed through the electric field (yellow), it was
deflected.
Thomson proposed that electrons were located throughout an atom like plums in
a pudding, as shown in this model
Rutherford model
The gold foil experiment
Top: Expected results: alpha particles passing through the plum pudding model of the atom
with negligible deflection.
Bottom: Observed results: a small portion of the particles were deflected by the
concentrated positive charge of the nucleus.
Thomson's plum pudding model was disproved in 1909 by one of his former
students, Ernest Rutherford, who discovered that most of the mass and positive charge of an
atom is concentrated in a very small fraction of its volume, which he assumed to be at the
very center.
In the gold foil experiment, Hans Geiger and Ernest Marsden (colleagues of Rutherford
working at his behest) shot alpha particles at a thin sheet of gold, measuring their deflection
with a fluorescent screen. Given the very small mass of the electrons, the
high momentum of the alpha particles and the nconcentrated distribution of positive charge
of the plum pudding model, the experimenters expected all the alpha particles to pass
through the gold sheet without significant deflection. To their astonishment, a small fraction
of the alpha particles experienced heavy deflection.
This led Rutherford to propose a planetary model in which a cloud of electrons surrounded
a small, compact nucleus of positive charge. Only such a concentration of charge could
produce the electric field strong enough to cause the heavy deflection.
Rutherford’s model of the atom had electrons surrounding the nucleus at a
distance. (This model does not show the true scale of sizes and distances.)
Bohr model
The planetary model of the atom had two significant shortcomings. The first is
that, unlike planets orbiting a sun, electrons are charged particles. An
accelerating electric charge is known to emit electromagnetic waves according
to the Larmor formula in classical electromagnetism; an orbiting charge should
steadily lose energy and spiral toward the nucleus, colliding with it in a small
fraction of a second. The second problem was that the planetary model could not
explain the highly peaked emission and absorption spectra of atoms that were
observed.
The Bohr model of the atom
Quantum theory revolutionized physics at the beginning of the 20th century,
when Max Planck and Albert Einsteinpostulated that light energy is emitted or
absorbed in discrete amounts known as quanta (singular, quantum). In
1913, Niels Bohr incorporated this idea into his Bohr model of the atom, in
which an electron could only orbit the nucleus in particular circular orbits with
fixed angular momentum and energy, its distance from the nucleus (i.e., their
radii) being proportional to its energy. Under this model an electron could not
spiral into the nucleus because it could not lose energy in a continuous manner;
instead, it could only make instantaneous "quantum leaps" between the
fixed energy levels. When this occurred, light was emitted or absorbed at a
frequency proportional to the change in energy (hence the absorption and
emission of light in discrete spectra).
Bohr's model was not perfect. It could only predict the spectral lines of
hydrogen; it couldn't predict those of multielectron atoms. Worse still,
as spectrographic technology improved, additional spectral lines in
hydrogen were observed which Bohr's model couldn't explain. In
1916,
Problems with the Bohr Model
-It violates the Heisenberg Uncertainty Principle because it considers electrons
to have both a known radius and orbit.
-The Bohr Model provides an incorrect value for the ground state orbital angular
momentum.
-It makes poor predictions regarding the spectra of larger atoms.
-It does not predict the relative intensities of spectral lines.
-The Bohr Model does not explain fine structure and hyperfine structure in
spectral lines.
-It does not explain the Zeeman Effect.
Arnold Sommerfeld (1868-1951) - was the German scientist added elliptical
orbits to the Bohr model , assumed that the orbits of electrons doesn't have to be
spherical but can also be elliptic. The electrons can move only on some, allowed
ellipses. He coined a second l number which was called the secondary quantum
number. The number defined the shape, the oblateness of an orbit. For n=1 the
orbit can be only spherical (l=0), for n=2 there are two orbits of different shapes
(l=0 - the elliptic one, l=1 - the spherical one). For any n there are n kinds of
shapes of the orbits. The electrons moving on the two orbits of the same n
number but of different shape have a bit different energies. That explaines the
discovered structure of the spectral lines.
Another thing improving the Bohr atom was the discovery that the orbits don't
have to lay in the same plane. They can be oriented in space on some defined
directions. Their orientation is defined by the magnetic quantum number ml.
The electron circulating on an orbit causes the magnetic field. If the system is
also placed in the outer magnetic field than the orbit of the electron places itself
in a special way. That means that its position makes the direction and the sense
of the magnetic field created by the electron the same as in the outer field. To
deflect it to another position there should be some more energy given to the
system. Sommerfeld proved that there is only some defined number of the
possible orbit's positions. The number is equal 2*l+l. Each position in the
magnetic field is of a bit different energy. The m number can name the values of
1 to -1. The phenomena of taking in the magnetic field the orbits of different
energies for the same n is the explanation of the splitting of spectral lines - the
Zeeman effect.
Zeeman effect
A splitting of spectral lines when the light source being studied is placed in a
magnetic field. Discovered by P. Zeeman in 1896, the effect furnishes
information of prime importance in the analysis of spectra. Each kind of spectral
term has its characteristic mode of splitting, and the types of terms are most
definitely identified by this property, Furthermore, the effect allows an
evaluation of the ratio of charge to mass of the electron and an evaluation of its
precise magnetic moment.
The normal Zeeman effect is a splitting into two or three lines, depending on the
direction of observation, as shown in the illustration. The light of these
components is polarized in ways indicated in the illustration. The normal effect
is observed for all lines belonging to singlet systems, those for which the spin
quantum number S = 0. The change of frequency of the shifted components can
be evaluated on classical electromagnetic principles.
Triplet observed in normal Zeeman effect
The anomalous Zeeman effect is a more complicated type of line splitting, so
named because it did not agree with the predictions of classical theory. It occurs
for any spectral line arising from a combination of terms of multiplicity greater
than one. Since multiplicity in spectral lines is caused by the presence of a
resultant spin vector S of the electrons, the anomalous effect must be attributed
to a nonclassical magnetic behavior of the electron spin.
The quadratic Zeeman effect, which depends on the square of the field strength,
is of two kinds. The first results from second-order terms, and the second from
the diamagnetic reaction of the electron when revolving in large orbits.
The inverse Zeeman effect is the Zeeman effect of absorption lines. It is closely
related to the Faraday effect, the rotation of plane-polarized light by matter
situated in a magnetic field.
Beside the described facts the one more was discovered - that the spectral line
consist of the two lines placed very close to each other - much closer than for
the two shapes of orbits. The fact couldn't be explained by the Bohr-Sommerfeld
model.
Uhlenbeck and Goudsmit explained the phenomena. The noticed that the
electron circulates not only around the nucleus but also rotates. It can rotate in
the two directions creating the rotary current flowing in the direction of the
rotation. The current induces the magnetic field which is directed with the field
created by the electron moving on the orbit or oppositely to it. So the fields sum
or substract. The electron causing the oppositely directed field is of a bit smaller
direction than the one causing the with directed field. So the spectral lines split.
The spin, magnetic quantum number is used for defining the direction of
rotation. The spin can be equal +1/2 or -1/2. That causes the noticed split of the
spectral
line
on
the
two
close
placed
ones.
After being improved the Bohr's theory very well described the spectrums of
hydrogen. The main lines are caused by the electron taking different orbits
(defined by the n). The lines consist of some very close placed lines for the
different shapes of orbits (defined by the l). In the magnetic field the lines
undergo a split because of the orbits taking defined planes in space. The last
thing is that the electron can rotate in two directions what causes the split of the
spectral lines on the two more ones. hydrogen - like atoms (f. e. He+). It
described the construction of atoms and of the orbits. So it was found a success.
But for some scientists its assumptions looked artificial.
In the case of a helium atom with two electrons in the 1s orbital, the Pauli
Exclusion Principle requires that the two electrons differ in the value of one
quantum number. Their values of n, l, and ml are the same; moreover, they have
the same spin, s = 1⁄2. Accordingly they must differ in the value of ms, which
can have the value of +1⁄2 for one electron and −1⁄2 for the other.
It is the underlying structure and symmetry of atomic orbitals, and the way that
electrons fill them, that determines the organisation of the periodic table and the
structure and strength of chemical bonds between atoms.