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Medical Chemistry Lecture By : Asst. LectTariq Al Mgheer of Medicine- Babylon University
PERIODIC TABLE
Chemists of the nineteenth century noticed that many elements have similar chemical
properties. Numerous attempts were made to arrange the elements into some systematic
fashion to emphasize these similarities. The most successful arrangement was arrived at
independently by two chemists, Dmitri Mendeleev and Lothar Meyer. They grouped the
elements with similar properties together into families and designed the first periodic table of
the elements. The modern periodic table is a direct descendant of the original MeyerMendeleev arrangement of the elements.
The periodic table consists of a number of columns called groups. The horizontal
rows are called periods. The elements in each gcouo have similar chemical and physical
properties. For example, the elements at the far right of the periodic table, called the noble
gases, are the least reactive of all the elements. The elements in group IA are called the
alkali metals. Those in group IIA are called the alkaline earth metals. The transition metals
are the elements located in groups IIIB to IIB. The halogens are the elements in group VIIA.
All the members of a group have similar physical and chemical properties.
The elements are divided into three classes based on their physical properties.
Elements that show a metallic luster when polished, are capable of being drawn out into
wire, can be hammered into sheets, and are good conductors of heat and electricity ars
classed as metals. Elements that do not have these properties are classed as nonmetals. A
class between these two is called the metalloids, or borderline elements. Elements of the
three classes are shown in the periodic table in Figure 1. There is a zigzag line on the right
side of the periodic table that separates the nonmetals on the right from the metals on the
left. Near this line are the metalloids. These elements such as silicon (Si) and germanium
(Ge), have some properties that are similar to those of nonmetals and some that are similar
to those of metals.
Only 90 of the 106 elements are found in nature. The others are prepared in the
laboratory by instruments and techniques. A very few elements, such as gold, silver, copper,
and platinum, exist naturally in a pure state. The rest are found in nature as parts of
compounds; that is, their atoms are combined with other-atoms to form compounds. Of the
90 elements found in nature, only eight make up 98 percent of all the compounds on earth.
These elements and their percentages are given in
Medical Chemistry Lecture By : Asst. Lect. Tariq Al Mgheer- College of Medicine- Babylon
UrUvsraitv
Table 1. Notice that the elements carbon, hydrogen, and nitrogen do not
appear in this table. These elements make up the compounds found in living
systems, yet each one accounts for less than 1 percent of all the elements on earth.
Table 1 The Most Abundant Elements on earth.
Element
Silicon
25.7
m
Weight
Percent
Calciu
3A———
—————
Sodiu
2.6
Aluminum
7.5
m
Potas
2.4
4.7
sium
Magn
1.9
Oxygen
Iron
Weight
Percent
49.5
Elem
ent
esium
Periodic Table Of The Elements
Most of the natural elements had been discovered by the end of the nineteenth
century. Furthermore, the idea that the elements are made up of atoms wis completely
accepted. Beginning in 1895, however, a number of subatomic particles were discovered
Their discovery overturned the easily understood world of indestructible atoms and laid the
foundation for our modern ideas about atoms and chemical reactions.
PARTS OF ATOMS
New laboratory tools and techniques have led to the discovery of many new
Medical Chemistry Lecture By : Asst. Lect. Tariq-H-Almgheer College of Medicine- Babylon University
subatomic particles during the past century. We need to consider only three of these
particles. With these three, we can constrict a model of the atom that satisfactorily accounts
for the chemical properties of all 106 atoms. These three particles are the electron, the
proton, and the neutron. Each of these particles has characteristic properties.
The electron is the particle responsible for electric current. It is a negatively charged
particle, which is defined as having one unit of negative charge. The symbol e is used to
represent an electron. The superscript minus sign indicates its negative charge.
The proton is a positively charged particle. Its charge is equal to but opposite that of
the electron. Therefore, a proton has one unit of positive charge. The symbol I-T represents
a proton. (Other symbols such is p and p"1" are sometimes used to represent a proton).The
plus sign placed as a superscript indicates its positive charge.
The neutron is a particle that has no charge. For this reason, it escaped detection until1932.
The symbol n represents a neutron.
All three of these particles have mass. But their masses are so small that it is difficult
to compare them with familiar objects. For example, the mass of a proton is 1.672 X 10"74 g.
Such masses are so small and so inconvenient to use that a new unit, the atomic mass unit
(amu), has been adopted. In this unit, a proton has a mass of 1,007 amu. .The mass of a
neutron is slightly less, 1.004 amu. The mass of an electron is still less, only (1 / 1837) the
mass of a proton. The characteristics of these three particles are summarized in Table 2.
Table 2 Three Important Subatomic Particles
NAME
CHARGE
MASS (AMU)
SYMBOL
Electron
-1
1/1837
e~
Proton
+1
1,007
H', P', p
Neutron
0
1.004
n
PUTTING THE PARTS TOGETHER
All atoms have two things in common.
First, they all have a nucleus which contains the neutrons and protons.
Second, their electrons occupy the space outside the nucleus. But here the similarity ends.
The differences in atoms of different elements are due to the fact that the atom have
differences numbers of protons in their nuclei. For example, the nucleus
Medical Chemistry Lecture By : Asst. Lect. Tariq-H-Almgheer- College of Medicine- Babylon University
of the helium atom contains two protons, whereas the nucleus of an oxygen atom contains
eight protons. The nucleus of an atom has a positive charge because the nucleus contains
protons. The charge of a nucleus is equal to the number or protons it contains. Thus, the
helium nucleus has a charge of +2, whereas the oxygen nucleus has a charge of +8.
An atom is electrically neutral. Therefore, the charge created by the protons in the
nucleus must be balanced exactly by an equal number of electrons outside the nucleus. For
example, the helium atom has two protons in its nucleus. To balance the charge of +2
created by these protons, it must have two electrons outside the nucleus. Similarly, an
oxygen atom has eight protons in its nucleus and it must have eight electrons outside the
nucleus.
Remember that neutrons, protons, and electrons are all very small particles. The
atom that contains these particles is also very small. A uranium atom, one of the largest,
has a diameter of 2.8 X 10'8 cm. This is so small that a length of 1 cm corresponds to 36
million uranium atoms placed side to side!
Atoms are small, but the nucleus is even smaller; it occupies only a small part of the
total volume of the atom. If one could magnify the size of an atom so that its nucleus was the
size of a baseball the atom would be a sphere approximately 16 km in diameter. Two
important points are made by this picture of an atom. First, the most massive particles of the
atom are concentrated in a very small volume of space. Second, most of the atom is empty
space.
The number of protons in the nucleus is an important property of an atom. This
number determines not only the number of electrons outside the nucleus, but also the atomic
number.
ATOMIC NUMBERS
At the turn of the century it was discovered that the number of positive charges on the
nuclei increases by one from atom to atom as one moves from one group to the next in any
period of the periodic table. For example, the positive charge on the hydrogen nucleus is +1;
on helium, +2; on lithium, +3. The positive charge is due to the protons in the nucleus.
Therefore, each element has one more proton in its nucleus than the element just before it in
the periodic table. This fact allows us to give numbers, called atomic numbers, to each
element. The atomic number of an element is equal to number of proton in its nucleus.
The atomic number of
Medical Chemists ,-ecture By : Asst. Lect. Tariq-H-Almgheer- College of Medicine- Babylon University
hydrogen is 1; of helium 2; of lithium 3. The atomic number of each element is given in the
periodic table (it is the number placed directly below the symbol of the element periodic
table).
All the atoms of an element have the same atomic number. The importance of this
statement became clear with the discovery of isotopes. ISOTOPES
Atoms of the same element can have different numbers of neutrons in their nuclei.
For example, most hydrogen atoms contain only one proton in their nuclei. A few hydrogen
atoms called deuterium atoms have nuclei that contain one proton and one neutron. The
nuclei of still other hydrogen atoms, called tritium atoms, have one proton and two neutrons.
All are atoms of hydrogen because their nuclei contain one proton (all have an atomic
number of one). However, the atoms differ in the number of neutrons in their nuclei. Atoms
whose nuclei have the same number of protons but different numbers of neutrons are called
isotopes. Hydrogen is not the only element that has isotopes. Many other elements have
two or more isotopes. Isotopes of an element have the same number of protons, but different
numbers of neutrons. Consequently, isotopes have different mass numbers and relative
atomic masses.
Mass Number Relative Atomic Masses
The mass number of an atom is the sum of the number of protons and neutrons
in its nucleus. The mass number of hydrogen is 1; that of deuterium is 2. The mass number
identifies a particular isotope of an element. This information can be added to the symbol of
the element. Thus, a specific isotope of an element is designated by writing the symbol of the
element with the atomic mass number placed as a superscript to the left. Usually, the atomic
number is also added as a subscript to the left. In this way we distinguish between isotopes
of the same element The three isotopes of hydrogen are designated as follows:
Mass Number(number of P &N)
1
2
1H
3
1H
1H
Symbol of elements
Atomic Number( number of P)
Isotope of carbon oxygen and sulfur are given in Table 4.
Symbol
Number of P
Number of n
Atomic number
Mass number
12
6C
6
6
6
12
13
6C
6
7
6
13
8
O16
8
8
8
16
17
8
9
8
17
32
16S
16
6
16
32
34
16S
16
18
16
34
36
16S
16
20
16
36
8O
Isotopes differ in mass because (they have different numbers of neutrons in their
nuclei. But what about the masses of different elements? Are they also different? Atoms of
different elements differ in their numbers of protons and neutrons. Therefore, it seems logical
to suppose that they do have different masses. The masses of different atoms are
determined by comparing them to a standard mass.
For atoms, the Standard mass, adopted by international agreement in 1961, is one
isotope of carbon
6
C12 called carbon 12. One atomic unit (amu) is defined as one-
twelfth the mass of a carbon -12 atom. This arbitrarily sets the atomic mass of 6 C12
as exactly 12 amu. The mass of other atoms are obtained by comparing them to
carbon 12. This is done experimentally by means of an instrument called a mass
spectrograph. The masses of all isotopes relative to carbon 12 can be determined by using
this instrument. Values of the masses of selected isotopes are given in Table5. Table 5
Isotooic Masses and Average Atomic Masses
M
ass
Ele
ment
I
sotope
Isotop
e Mass
N
Nat
ural
Abundance
(amu)
(%)
Average
Atomic Mass of
Elements
(amu)
umber
Nitr
14.0067
ogen
1
7
N14
4
14,00
3
99.6
3
1
7
N15
5
15.00
0.37
0
Bor
10.811
on
1
5
B10
0
5
11
B
10.01
3
1
1
19.6
0
11.00
9
80.4
0
Sulf
32.064
ur
3
1
32
6S
2
31.97
2
95.0
0
Medical Chemistry Lecture By : Asst. Lect. Tariq-H-Almgheer- College of Medicine- Babylon University
16
33
32.971
0.76
16
34
33.967
4.22
16
36
35967
0.014
12
24
23.985
78.70
12
25
24.986
10.13
26
25.983
11.17
S33
S34
S36
Magn
24.312
Mg24
esium
Mg25
26
12Mg
Notice that the atomic mass of an isotope is very close to, but not exactly equal to, a
whole number. In fact. the relative mass of an isotope is almost the same as its mass
number. What about the atomic masses of the elements? This can be seen from the values
of the atomic masses of several elements given in Table 5. The reason for this is that an
element is made up of a number of isotopes. Therefore, the mass of an element is a
weighted average of the masses of all its isotopes. For example, it is Known that carbon is
made up of 98.89 percent
6C!3 and 1.11percent 6C13 the -----------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------
average mass of carbon is the weighted average of the masses of these two
isotopes:
Average atomic mass of C= (0.9889) (12.00)+(0.0111) (13.00)= 12.01 This
is the way all the atomic masses of the elements are obtained.
The terms mass and weight are often used interchangeably. As a result, the atomic
mass of an element is also called its atomic weight.
The composition of the nucleus determines many properties of an atom. The number
of protons in the nucleus determine the atomic number and the number of electrons outside
the nucleus. Isotopes are atoms that have the same number of protons but different numbers
of neutrons in their nuclei. Finally, the mass number of an atom is determined by the number
of neutrons and protons in its nucleus.
ELECTRON ARRANGEMENT
The electron occupy most of the space in an atom, but they do not have unlimited
freedom to go anywhere within the atom. Electrons are located in certain well-defined shells
in the space around the nucleus.
The various shells in which electrons are located are given numbers. Tl-ie first is
shell number 1 (also sometimes called the K shell). This shell is nearest the nucleus. An
electron in this shell has the lowest energy of any electron in the atom.
7
IVIedical Chemistry Lecture By : Asst. Lect.Tariq-H-Almgheer- College of Medicine- Babylon University
3- The energy levels of the shells are closer together the farther the shells are from the
nucleus. In fact, there is an overlap of the energies of the atomic orbitals in shells 3 and 4. As
a result, the 4s atomic orbital is lower in energy than the 3d orbitals.
The concept of atomic orbitals and their relative energy levels is extremely important to
our understanding of the arrangement of electrons in atoms. Using the relative energy levels
of these orbitals (Figure2) and the rule that each atomic orbital can contain a maximum of
two electrons, we can describe the arrangement of electrons in any atom. The electron
arrangements of the first 20 elements are given in Table 7. Table 7 Electron Arrangement
of the First Twenty Elements
Element
Atomic Orbitals
Is
2s
2px
2Py
2Pz
He
Li
XX
x
Be
B
C
N
0
XX
XX
XX
XX
XX
F
Ne
Na
3s
3px
Spy
3Pz
4s
XX
XX
XX
XX
XX
x
x
x
XX
x
x
x
x
x
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
x
XX
XX
x
Mg
XX
XX
XX
XX
XX
XX
Al
XX
XX
XX
XX
XX
XX
x
Si
XX
XX
XX
XX
XX
XX
x
x
P
S
Cl
Ar
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
x
XX
XX
XX
x
x
XX
XX
x
x
x
XX
K
XX
XX
XX
XX
XX
XX
XX
XX
x
Ca
XX
XX
XX
XX
XX
XX
XX
XX
XX
H
Several important points should be noted about the electron arrangements of the
elements given in Table 7.
First. whei.Jver there is more than one atomic orbital of the same energy, one
electron occupies each atomic orbital of equal energy first before two electrons are placed in
any one atomic orbital. For example, a carbon atom has two electrons in the 1s atomic
orbital, two in the 2s atomic orbital, and one electron in each of two different 2p atomic
orbitals.
Second, the 4s atomic orbital is lower in energy than the 3d atomic orbitals.
IVIedical Chemistry Lecture By : Asst. Lect.Tariq-H-Almgheer- College of Medicine- Babylon University
3- The energy levels of the shells are closer together the farther the shells are from the
nucleus. In fact, there is an overlap of the energies of the atomic orbitals in shells 3 and 4. As
a result, the 4s atomic orbital is lower in energy than the 3d orbitals.
The concept of atomic orbitals and their relative energy levels is extremely important to
our understanding of the arrangement of electrons in atoms. Using the relative energy levels
of these orbitals (Figure2) and the rule that each atomic orbital can contain a maximum of
two electrons, we can describe the arrangement of electrons in any atom. The electron
arrangements of the first 20 elements are given in Table 7. Table 7 Electron Arrangement
of the First Twenty Elements
Element
Atomic Orbitals
Is
2s
2px
2Py
2Pz
He
Li
XX
x
Be
B
C
N
0
XX
XX
XX
XX
XX
F
Ne
Na
3s
3px
Spy
3Pz
4s
XX
XX
XX
XX
XX
x
x
x
XX
x
x
x
x
x
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
x
XX
XX
x
Mg
XX
XX
XX
XX
XX
XX
Al
XX
XX
XX
XX
XX
XX
x
Si
XX
XX
XX
XX
XX
XX
x
x
P
S
Cl
Ar
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
XX
x
XX
XX
XX
x
x
XX
XX
x
x
x
XX
K
XX
XX
XX
XX
XX
XX
XX
XX
x
Ca
XX
XX
XX
XX
XX
XX
XX
XX
XX
H
Several important points should be noted about the electron arrangements of the
elements given in Table 7.
First. whei.Jver there is more than one atomic orbital of the same energy, one
electron occupies each atomic orbital of equal energy first before two electrons are placed in
any one atomic orbital. For example, a carbon atom has two electrons in the 1s atomic
orbital, two in the 2s atomic orbital, and one electron in each of two different 2p atomic
orbitals.
Second, the 4s atomic orbital is lower in energy than the 3d atomic orbitals.
Medical Chemistry Lecture By : Asst. Lect. Tariq-H-Almgheer- College of Medicine- Babylon University
energy needed to remove the first electron from an atom is called its first ionization
energy. The lower the ionization energy, the easier it is to remove an electron from the atom.
A graph of the ionization energies of the atoms and their atomic numbers reveals a periodic
relationship, as shown in Figure 3.
Figure 3 The first ionization energies of the first 20 elements plotted against their
atomic numbers.
Starting with, lithium the values of the first ionization energy of the elements of the
first period generally increase to reach a maximum at neon. Then there is a sudden
decrease at sodium The values then begin to increase again to reach another maximum at
argon. From this graph, we can conclude that the alkali metals (lithium, sodium, and
potassium) lose an electron most easily. As we proceed along a period, it becomes
increasingly difficult for an element to lose an electron. The noble cases are the most
reluctant of all the elements to lose an electron.
Electron affinity is the energy released when an atom gains an electron. It is a
measure of the ability of an atom to attract an electron. The higher the value of the electron
affinity, the greater the ability of an atom to gain an electron. The electron affinities of the first
20 elements are plotted against their atomic numbers in Figure 4.
Figure 4 The electron affinity of the first 20 elements plotted against their atomic
numbers.
A periodic relationship is apparent in Figure 4.Starting with helium, the values of the
eleciron affinities of the elements in the first period increase and decrease slightly until
oxygen and fluorine are reached. These two elements have the highest electron affinities in
the first period. Then there is a sudden decrease at neon. The values in the second period
parallel those in the first row. The values of the elements in group VIA and VIIA (sulfur and
chlorine, respectively) are again the highest in the period. As before, there is a sudden
decrease when argon is reached. From this graph, we conclude that the noble gases are the
most reluctant to accept an electron, and the halogens are the most willing of all elements to
accept an electron.
There is an important relationship between the electron arrangement of an atoms and
the value of their ionization energies and electron affinities. The noble gases have high
ionization energies and low electron affinities. This means that the atoms of these elements
are reluctant to lose or to gain an electron. The electron arrangement of the noble gases is
particularly stable. This arrangement is so stable that many other atoms either lose or gain
electrons to achieve this arrangement. For example, a sodium atom easily loses an electron
to form a positively charged sodium ion. A sodium ion has 11 protons but only 10 electrons.
Therefore, it has a net positive charge of +1. The electron arrangement of sodium ion (Na^)
is identical to that of the noble gas neon. This is the reason the alkali metals have such low
ionization energies. They easily lose an electron to form the electron arrangement of the
preceding noble gas.
Medical Chemistry Lecture By : Asst. Lect. Tariq-H-Almgheer-College of Medicine- Babylon University
Electron affinity and ionization potential are measures of two different processes. One
is the loss of an electron from an atom, and the other is a gain of an electron. To indicate
which of these two processes is dominant in an atom, the idea of electronegativity was
developed. Electronegativity is the tendency of an atom to attract electron to it. The
electronegativity of an atom was originally calculated as the average of its electron affinity
and its ionization potential. Since then, other scales of electronegativity have been proposed.
The scale in most common use today is the one proposed by Linus Pauling. It assigns a
value of 4 to fluorine and compares all other atoms to this standard value. For example,
sodium, an atom with a relatively small attraction for electrons, has an electronegativity value
of only 0.9. Carbon, whose electron attraction is moderate, has a value of 2.5, The higher the
value of the electronegativity of an atom, the greater the tendency to attraction electrons to
itself.
A graph of the Pauling electronegativity values of the atoms and their atomic
numbers has a periodic relationship, as shown in Figure 5.
Figure 5 The electronegativities of the first 20 elements plotted against their atomic
numbers.
Starring with lithium, the values of electronegativity of the first period increase to
reach a maximum at fluorine. Then there is a sudden decrease at sodium, followed by an
increase until another maximum is reached at chlorine. In general, there is a decrease in the
electronegativity with fall within a particular group. For example, the electronegativity of the
halogens decrease in the order fluorine, chlorine, bromine and iodine.
The importance of electronegativity is that we can tell the kind of bond formed
between two atoms simply by looking at the difference in their electronegativities.'