ATOMIC STRUCTURE
Atom is the particle that could not be divided any further and extremely small,
invisible, and indivisible.
Carbon and oxygen can react to form carbon monoxide, In the reaction, one
carbon atom (C) can also combine with two atoms of an oxygen molecule
(O2) to form a molecule of carbon dioxide CO2.
1 atom of C
+
1 molecule of O2
1 molecule of CO2
Rutherford described the atom as having a central positive nucleus. The
entire mass of the atom is concentrated in its nucleus and the rest of the
atom was mostly empty space. The observation that, the mass of an electron
is negligible compared to the mass of a proton or a neutron, indicated that
the protons and the neutrons are located in the nucleus, while the electrons
are found in the outer regions of the atom.
The positive charge of the
nucleus is determined by the number of the protons it contains. As protons
and electrons have equal but opposite charges, it follows that in an
electrically neutral atom the number of protons must be the same as the
number of electrons.
Rutherford proposed that the electrons (located in the outer regions of the
atom) orbit the nucleus in the same manner that the Earth and other planets
orbit the sun. Because of this analogy with our planetary system, Rutherford's
model is often referred to as the solar-system model of the atom. The
model makes no assumptions about the distance of the electrons from the
nucleus.
Bohr assumes that the electrons orbit around the nucleus on the surfaces of
imaginary spherical shells (levels). These electron shells are concentric about the nucleus in the same
way as the successive layers of an onion are packed together.
Bohr model of the atom is also
known as the electron shell model,
Internal Structure of the Atom
Atom is made up of three major subatomic particles ; protons, neutrons, and
electrons.
Proton
The proton is an elementary particle with a mass of 1.67 × 10-24 g and has
the smallest unit of positive charge. According to the fundamental laws of
electricity, protons will repel each other, attract particles with negative
charges, and do not interact with particles that carry no charge.
Electron
The electron has the lowest mass, only 1/1836 that of a proton and has a
negative charge which is equal in magnitude to that of the proton. Thus,
electrons repel each other, attract protons, and do not interact with
electrically neutral particles.
Neutron
The neutron is a subatomic particle with a mass almost equal to that of the
proton but with no electrical charge. Because of its electrically neutral nature,
this particle will neither attract nor repel positively charged protons,
negatively charged electrons, or other neutrons.
Characteristics of electrons and nucleons:
Particle
Charge Mass (amu)*
Mass (kg)
Relative mass
Electron
-1
0.000549
0.9108 x 10 -30
1
Proton
+1
1.00728
1.6721 x 10 -27
1
0
1.00867
1.6744 x 10 -27
1/1836
Neutron
* amu = 1 atomic mass unit = 1.66 x 10
-27
kg = 1 / 12 of the mass of
12C.
Electron Shells
The seven electron shells are labeled with integers n = 1, 2, 3, 4, 5, 6 and 7 starting from the shell closest to the
nucleus. Another convention for labeling the electron shell in by means capital letters; K (n = 1), L (n = 2), M (n =
3), N (n = 4) and so forth. An electron in a shell with a relatively low value of n is at a shorter distance from the
nucleus than an electron in a shell with a higher value of n. Since the principle quantum number (n) is a measure of
the distance of an electron from the nucleus, it is also a measure of the energy possessed by that electron. Electrons
in shells with low value n have a lower energy than electrons in shells with higher value of n.
Bohr deduced that electrons inside an atom possess different energies where
electrons in the
first
orbit belong to the
orbit belong to the
first energy level,
second energy level…..etc…..
electrons in the
second
Each energy level of an atom
could only accommodate a certain number of electrons. The maximum
number of electrons that can populate a certain energy level is given by the
following formula.
Maximum number of electrons in an energy level = 2 n2
where: n is the specific energy level ( n= 1, 2, 3, 4, ……………… 7 )
n = 4 (N)
n = 3 (M)
n = 2 (L)
n = 1 (K)
2e8e18e32e-
The electron shell model showing
the maxium number of electrons (e-)
that can be accommodated in each shell
Atomic Number and Nucleon Number:
The nucleus of an atom always contains a whole number of protons, exactly
equal the number of electrons in the neutral atom.

Atomic number is known as the number of protons in the nucleus of
an atom.
Atomic number = Number of protons = Number of electrons

Nucleon (Mass) number is known as the sum of the numbers of
protons and neutrons.
Nucleon number = Number of protons + Number of neutrons

Number of neutrons = Nucleon number – Atomic number
Protons and neutrons, collectively called nucleons. The number of protons is
called the atomic number of the element and denoted by Z. The number of
neutrons is denoted by N, and the sum of the protons and neutrons, Z + N,
is called the mass number, denoted by A.
The symbolic representation of an element, X is given by
A
ZXN
.
For
example, sodium has 11 protons and 12 neutrons with a total of 23 nucleons.
Thus, it is represented as
23
11Na12
It is observed that atomic nuclei containing an odd number of protons or
neutrons are normally less stable than those with an even number of protons
or neutrons. Thus, nuclei with even numbers of protons and neutrons are
more stable, whereas those with odd numbers of protons and neutrons are
less stable.
There are about 270 stable atoms of normally occurring elements.
The
stability of these elements is dictated by the configuration of protons and
neutrons. The ratio of the number of neutrons to the number of protons
(N/Z) is an approximate indicator of the stability of a nucleus. The N/Z ratio
is 1 in low-Z elements such as
12C
6
,
14N
7
,
and
16O
8
, but it increases with
increasing atomic number of elements. For example, it is 1.40 for
1.54 for
208Pb
82
127I
53
and
.
Nuclear binding energy:
According to the classical electrostatic theory, the nucleus of an atom cannot
exist as a single entity, because of the electrostatic repulsive force among the
protons in the nucleus.
The stability of the nucleus is explained by the
existence of a strong binding force called the nuclear force, which overcomes
the repulsive force of the protons.
The mass of a nucleus is always less than the combined masses of the
nucleons in the nucleus. The difference in mass is termed the mass defect,
which has been used as binding energy for all nucleons in the nucleus.
The average binding energy of a nucleon is equal to the total binding energy
(calculated from the mass defect) divided by the number of nucleons. It is of
the order of 6 – 9 MeV.
Note that whereas the binding energy of the
nucleons is in the megaelectron volt (MeV) range, the electron binding energy
in the atomic orbital is of the order of kiloelectron volts (keV), a factor of
1000 lower.
Nuclear nomenclature:
Nuclide:
It is an atomic species with a definite number of protons and
neutrons arranged in a definite order in the nucleus.
Radionuclides: They are those nuclides that are unstable and thus decay
by emission of particles or electromagnetic radiations or by spontaneous
fission.
Isotopes: They
are the nuclides having the same atomic number Z but
different mass number. Isotopes exhibit the same chemical properties.
Examples of carbon isotopes are
11
6C,
12
6C,
and
13
6C.
The most common isotope of the element oxygen has 8 protons and 8
neutrons in the nucleus of one of its atoms. The atomic number of this
isotope of oxygen is, therefore 8 and the nucleon number is 16. The great
majority (99.759%) of oxygen atoms in the nature occur as this isotope.
Hydrogen has three isotopes. The common isotope has a nucleus that
contains one proton only. The second one exist in every 5000 hydrogen
atoms has a nucleus that contains one proton and one neutron. This latter
isotope has twice the mass of an ordinary hydrogen atom and is called heavy
hydrogen or deuterium (D). An even smaller number of hydrogen atoms, 1 in
1017, has a nucleus with one proton and two neutrons. This isotope is called
super heavy hydrogen or tritium (T).
2
1
Hydrogen has three isotopes:
H ,
1
3
H
1
and
2
1
H ,
1
1
3
D
1
and
1
H
T
Isotones: They are the nuclides having the same number of neutrons N but
different numbers of protons. Examples of isotones are:
132
53I,
134
55Cs,
133
54Xe,
and
each having 79 neutrons.
Isobars: They are the nuclides with the same number of nucleons: that is,
the same mass number A, but a different combination of protons and
neutrons. For example:
82Y,
82Sr,
82Rb,
and
82Kr
are all isobars having the
mass number 82.
Isomers:
They are the nuclides with the same number of protons and
neutrons, but having different energy states and spins.
99Tc
and
99mTc
are
isomers of the same nuclide. Individual nuclides can exist in different energy
states above the ground state due to excitation. These excited states are
called the isomeric states, which can have a lifetime varying from picoseconds
to years. When the isomeric states are long-lived, they are reffered to as
metastable states. These states are donated by m as in
99mTc.
Electron Configuration of the Elements
The arrangement of electrons in an atom is called the electron
configuration. When electron fill the energy levels, it fills the lowest energy
level first.
Example: For a hydrogen atom, H, has
an atomic number 1,
the one electron goes
into the first energy level, the K shell (n = 1).
Example: For a lithium atom, Li, has an atomic number 3, two of the three
electrons go into the first energy level (K shell) while the third electron goes
into the second energy level (L shell). This electron in the outer energy level
is called the valence electron. The two electrons in the first energy level are
called the core electrons.
Problem: Give the electron configuration for silicon (atomic number 14).
Silicon, Si, atomic number 14 and hence 14 electrons. The first shell (K shell)
can accommodate 2 electrons, and the second shell (L shell) can hold 8
electrons. That leaves 4 electrons to be accommodated in the third shell (M
shell).
n = 3 (M)
n = 1 (K)
n = 2 (L)
n = 1 (K)
n = 1 (K)
14+
3+
1+
Hydrogen,
n = 2 (L)
1H
Lithium,
3Li
Silicon,
14Si
According to Bohr model of the atom
The Quantum Mechanical Model of the atom presents a more accurate model
of the atom. We will take a look at this model and summarize the results
based on these mathematical calculations without carrying them out
ourselves.
The Quantum Mechanical Model assumes that each shell is subdivided into
several number of sublevels (s, p, d and f ).
There is only one s-type orbital - There are three p-type orbitals,
There are five d-type orbitals - There are seven f-type orbitals
The first shell (K) contains only one orbital s,
The second shell (L) subdivided into two sublevels (s and p orbitals),
The third shell (M) subdivided into three sublevels (s, p and d orbitals), while
The fourth shell (N) and the other shells (n = 5, 6 and 7) subdivided into four
sublevels (s, p, d and f orbitals)
Each orbital can contain a maximum of two electrons. Wolfgang Pauli states that
if two electrons occupy the same orbital they must have opposite spin. This is
known as the
Pauli exclusion principle.
Summary: The distribution of electrons in each energy level
Energy
Level, n
Type of
Atomic Orbital
Number of
Atomic Orbitals
Maximum Number of
Electrons per Sublevel
1
1s
1
2
2
2s
1
2
2p
3
6
3s
1
2
3p
3
6
3d
5
10
4s
1
2
4p
3
6
4d
5
10
4f
7
14
3
4, 5, 6, 7
Maximum
Total Number
of Electrons
2
8
18
32
Numbers on the last column is equivalent to the prediction using the formula 2 n2
There is a way to represent precisely the electron arrangement in atoms. Let's
take a look at the simplest atom, hydrogen.
A hydrogen atom has 1 electron. That electron will occupy the lowest
principal energy level, n = 1, and the only sublevel, s. We denote the electron
configuration of hydrogen as 1s1.

Helium has 2 electrons; the 2 electrons both occupy the s sublevel
in principal energy level 1.
o

Helium's electron configuration is 1s2
Lithium has 3 electrons; 2 of the 3 electrons occupy the s sublevel
in principal energy level 1. The 3rd electron must go in the next
available sublevel, 2s.
o

Lithium's electron configuration is 1s2 2s1
Beryllium has 4 electrons; 2 of the 3 electrons occupy the s
sublevel in principal energy level 1. The 3rd and 4th electrons must go
in the next available sublevel, 2s. Beryllium's electron configuration is
1s2 2s2
4f
4d
4b
3d
4s
3p
3s
2p
2s
1s
The arrangement of sublevels in order of increasing energy.
Electron configuration of
11
5B,
1s2 2s2 2p1
11
5B
,
12
6C
,
and
14
7N
12
6C,
1s2 2s2 2p2
14
7N,
1s2 2s2 2p3
Electron configuration of
18Ar
,
20Ca
,
30Zn
and
18Ar,
1s2
2s2
2p6
3s2
3p6
20Ca,
1s2
2s2
2p6
3s2
3p6
4s2
30Zn,
1s2
2s2
2p6
3s2
3p6
4s2
3d10
36Kr,
1s2
2s2
2p6
3s2
3p6
4s2
3d10
36Kr
4p6
Often, to save space, electron configuration starts with the preceding nobel
gas ( 2He ,
10Ne, 18Ar, 36Kr, 54Xe
and
86Rn).
For example, in case of the elements sulfur = electron configuration of the
element Neon , [Ne] + 3s2 3p4. Howevere, in case of the element Nickel =
electron configuration of the element Argon , [Ar] + 4s2 3d8.
Element Electron configuration
Abbreviated electron configuration
16S
1s2 2s2 2p6 3s2 3p4
[Ne] 3s2 3p4
28Ni
1s2 2s2 2p6 3s2 3p6 4s2 3d8
[Ar] 4s2 3d8
Filling of sublevels and the periodic table:
The atoms of the group1 elements all have one s electron in the outermost
principal energy level. In each group2 atom, there are two s electrons in the
outermost principal energy level. A similar relationship applies to the elements
in any group:
" The atoms of elements in a group of the periodic table have the same
distribution of electrons in the outermost principal energy level." This means
that the order in which electron sublevels are filled is determined by position
in the periodic table.
Notice the following points:

The elements in group 1 and 2 are filling an s sublevel. Thus, Li
and Be in the second period fill the 2s sublevel. Na and Mg in the
third period fill the 3s sublevel, and so on.

The elements in group 13 through 18 (six elements in each period)
fill
p sublevels, which have the capacity of six electrons. In the
second period, the 2p sublevel starts to fill with B (Z=5) and is
completed with Ne (Z=10).
In the third period, the elements Al
(Z=13) through Ar (Z=18) fill the 3p sublevel.

The transition metals, in the center of the periodic table, fill d
sublevels.

The two sets of 14 elements called Lanthanides & Actinides listed
separately at the bottom of the table are filling f sublevels with a
principal quantum number two less than the period number.

14 elements in the sixth period (Z=57 to 70) are filling the 4f
sublevel.
These elements are sometimes called Rare earths or,
more commonly nowadays, Lanthanides, after the name of the first
element in the series, Lanthanum ( 57La).

14 elements in the seventh period (Z=89 to 102) are filling the 5f
sublevel.
The first element in this series is Actinium (
collectively, these elements are reffered to as Actinides.
elements are radioactive.
increasing atomic number.
89Ac),
All these
Their stability decreases rapidly with
The longest-lived isotop of Nobelium
(102No) has a half-life of about 3 min. Nobelium and the preceding
element, Mendeleevium (101Md), were identified in
containing one to three atoms of No or Md.
Electronic Structure and The Periodic Table
samples
The hydrogen atom:
The hydrogen atom, containing a single electron, has played a major role in
the development of models of electronic structure.
Quantum Numbers, Energy levels, and orbitals:
There are three quantum numbers, given the symbols n, ℓ, and mℓ. A wave
function corresponding to a particular set of three quantum numbers (e.g.,
n =2 ,
ℓ= 1,
mℓ = 0 ) is referred to as an atomic orbital.
Orbitals differ from one another in their energy and in the shape and spatial
orientation of their electron cloud. A fourth quantum number (ms) is required
to completely describe a specific electron in a multi-electron atom.
Each electron in an atom has a set of four quantum numbers, n, ℓ, mℓ, ms.
First quantum number (n), Principal energy levels:
This number, n, comes from the Bohr model of the hydrogen atom, where the
energy depends only upon n.
En = - RH / n2 ............................ (Bohr equation)
Where: En is the energy of the electron
RH is the quantity called the Rydberg constant (2.18 x 10-18 j)
n is an integer called the principal quantum number.
In another atoms, the energy of each electron depends mainly, but not
completely, upon the value of n. As n increases, the energy of the electron
increases, and, on the average, it is found farther out from the nucleus. The
quantum number (n) can take on only integral values, starting with 1:
n = 1, 2, 3, 4, .........
In an atom, the value of n designates what we call a principal energy level.
Thus, an electron for which n =1 is said to be in the first principal level. If
n=2, we are dealing with the second principal level, and so on.
Second quantum number (ℓ), Sublevels (s, p, d, f):
Each principal level includes one or more sublevels. The sublevels are denoted
by the second quantum number, l. The general shape of the electron cloud
associated with an orbital is determined by l. Large values of l produce more
complex shapes.
The quantum numbers, n, and ℓ are related; l can take on any integral value
starting with 0 and going up to a maximum of (n-1).
That is, ℓ = 0, 1,
2,......., (n-1).
If n = 1, there is only one possible value of ℓ, namely 0. This means that, in
the first principal level, there is only one sublevel, for which ℓ = 0. If n=2,
two values of ℓl are possible, 0 and 1.
In other words, there are two
sublevels (ℓl = 0 and ℓ = 1) within the second principal energy level.
Similarly,
If n = 3
ℓ = 0, 1, 2
(three sublevels)
If n = 4
ℓ = 0, 1, 2, 3
(four sublevels)
Another method is commonly used to designate sublevels. Instead of giving
the quantum number, ℓ, the letters s, p, d, or f indicate the sublevels ℓ=0, 1,
2, or 3, respectively. That is,
quantum number, ℓ,
0
1
2
3
Type of sublevel
s
p
d
f
Sublevels designations for the first four principal levels.
n
1
ℓ
0
Sublevels 1s
2
3
4
0
1
0
1
2
0
1
2
3
2s
2p
3s
3p
3d
4s
4p
4d
4f
Sublevels increase in energy in the order : ns < np < nd < nf
Thus, a 2p sublevel has a slightly higher energy than a 2s sublevel. By the
same token, when n = 3, the 3s sublevels has the lowest energy, the 3p is
intermediate, and the 3d has the highest energy.
Third quantum number (mℓ), Orbitals:
Each sublevel contains one or more orbitals, which differ from one another in
the value assigned to the third quantum number, mℓl. This quantum number
determines the direction in space of the electron cloud surrounding the
nucleus. The value of mℓ is related to that of ℓl. For a given value of ℓl, mℓ
can have any integral value, including 0, between ℓ and –ℓ; that is,
mℓ =ℓ, .........., +1, 0, -1,........., -ℓ
To illustrate how this rule works, consider an s sublevel (ℓ = 0). Here mℓ can
have only one value, 0. This means that an s sublevel contains only one
orbital, referred to as an s orbital.
The electron cloud associated with an s orbital is spherically symmetrical; the
density of the cloud varies with distance from the nucleus but is independent
of direction. Most commonly, an s orbital is shown as a simple sphere.
Z
y
The radius of the sphere indicates the region within which there is a specified
probability of finding the electron.
For a p sublevels (ℓ = 1),
mℓ =1, 0, or -1. Within a given p sublevels, there
are three different orbitals described by the quantum numbers mℓ = 1, 0,
and -1. Commonly, p orbitals are referred to as px, py, and pz orbitals.
px orbital
pz orbital
py orbital
The electron density in p orbitals (b) is concentrated along the x, y, or z axis.
The three p orbitals are directed at 90° angles to each other.
For the d and f sublevels:
d sublevels: ℓ = 2 mℓ =
2, 1, 0 , -1, -2
f sublevels: ℓ = 3 mℓ = 3, 2, 1, 0 , -1, -2, -3
5 orbitals
7 orbitals
In general, for a sublevel of quantum number ℓ, there are a total of 2ℓ+1
orbitals.
Fourth quantum number (ms), Electron spin:
The quantum number ,ms, is associated with the spin of the electron. An
electron has magnetic properties that correspond to those of a charged
particle spinning on its axis.
Either of two spins are possible, clockwise or counter clockwise.
N
N
N
S
S
S
S
Repulsion
N
some attraction
ms number is not related to n, ℓ, or mℓ. It can have either of two possible
values ms = +½ or - ½
Electrons that have the same value of ms (i.e., both +½ or both - ½ ) are
said to have parallel spins. Electrons that have different ms values (i.e., one
+½ and the other - ½ ) are said to have opposed spins.
Pauli Exclusion Principle:
This rule relates to the four quantum numbers that characterize an electron in
an atom. It requires that no two electrons in an atom can have the same set
of four quantum numbers.
The Pauli exclution principle requires that no more than two electrons can fit
into an orbital. Moreover, if two electrons occupy the same orbital they must
have opposed spins.
To see that this is the case, consider the 2s orbital. Any electron in this
orbital must have
n =2
ℓ=0
mℓ= 0
To satisfy the Pauli exclusion principle, the electrons in this orbital must have
different ms value. But there are only two possible values of ms. Hence,
only two electrons can enter the orbital. If the orbital is filled, one electron
must have ms = +½ and the other ms = -½ , the two electrons must have
opposed spins.
Capacities of principal levels, sublevels and orbitals:
1. Each principal level of quantum number, n , contains a total of n sublevels.
2. Each sublevel of quantum number , ℓ , contains a total of 2ℓ +1 orbitals;
that is, an S sublevel (ℓ = 0) contains 1 orbital
p sublevel (ℓ = 1) contains 3 orbitals
d sublevel (ℓ = 2) contains 5 orbitals
f sublevel (ℓ = 3) contains 7 orbitals
3. Each orbital can hold two electrons (2e-) , which must have opposed spins.
Allowed sets of quantum numbers for electrons in atoms:
Level n
1
Sublevel ℓ
0
0
Orbital mℓ
0
0
1
0
-1
0
1
0
-1
2
1
0
-1
-2
1s
2s
2px
2py
2pz
3s
3px
3py
3pz
3d
3d
3d
3d
3d
Spin ms
2
3
1
0
1
2
=+½
= -½
Example:
(a) What is the capacity for electrons of the 3d sublevels?.
(b) How many electrons can fit into the principal level for which n = 4.
Solution:
a) Each d sublevel contains five orbitals, so its capacity is 5 x 2 e- = 10 e-.
b) if n = 4 , there must be four sublevels, 4s, 4p, 4d, 4f.
S sublevel has one orbital, p sublevel has three orbitals, d sublevel has five
orbitals and f sublevel has seven orbitals.
Therefore, 1 (2 e-) + 3 (2 e-) +5 (2 e-) +7 (2 e-) = 32 e-
Capacities of electronic levels and sublevels in atoms:
Level n
Total № of electrons in level, 2 n2
Maximum № of electrons in sublevels,
2 (2ℓ +1)
s
p
d
F
1
2
2
-
-
-
2
8
2
6
-
-
3
18
2
6
10
-
4
32
2
6
10
14
Periodic trends in the properties of atoms:
Periodic law: The chemical and physical properties of elements are a
periodic function of atomic number. Atomic radius, ionic radius,
ionization energy and electronegativity vary horizontally and vertically in
the periodic table.
Atomic Radius:

Atomic radius is taken to be one half the distance of closest approach
between atoms in an elemental substance.
0.256 nm
Cu
Atomic radius = 0.256 / 2 = 0.128 nm

Atomic radii decrease across a period from left to right in the periodic
table:
Atom
Li
Radius 0.152
Be
B
C
0.111
0.088
0.077
N
0.070

Atomic radii increase down a group in the periodic table:
Li
Be
0.152
0.11
increase
Na
0.189
Mg
0.160
K
0.231
Ca
0.197
decrease
As effective nuclear charge increases, outer electrons are pulled in more
tightly, and atomic radius decreases.
Ionic Radius:
Negative ions are always larger in size than the atoms from which they are
derived, whereas positive ions are smaller. As a result of these effect, anions
in general are larger than cations.
Positive ion
(Cation)
Atom
Negative ion
(Anion)
The ionic radius increases moving down a group in the periodic table.
Moreover, the radii of both cations and anions decrease from left to right
across a period.
Li+
Be2+
0.060
0.031
Na+
Mg2+
0.095
0.065
Al3+
0.050
K+
Ca2+
Ga3+
0.133
0.099
0.62
The difference in radii between atoms and ions can be explained quite simply.
A cation is smaller than the corresponding metal atom between the excess of
protons in the ion draws the outer electrons in closer to the nucleus.
In
contrast, an extra electron in an anion adds to the repulsion between outer
electrons, making a negative ion larger than the corresponding nonmetal
atom.
O2-
F-
0.140
0.136
S2-
Cl-
0.189
0.181
Se2-
Br-
0.198
0.195
Te2-
I-
0.221
0.216
Ionization Energy:
Ionization energy is a measure of how difficult it is to remove an electron
from a gaseous atom.
Energy must always be absorbed to bring about
ionization, So ionization energy is always positive quantity.
The (first) ionization energy is the energy change for the removal of the
outermost electron from a gaseous atom to form a positive ion (+1).
M(g)
M(g)+ + e-
∆ E1 = first ionization energy.
The more difficult it is to remove electrons, the larger the ionization energy.
Ionization energies increase across the periodic table from left to
right.
Ionization energies decrease moving down the periodic table.
There are an inverse correlation between ionization energy and atomic radius.
The smaller the atom, the more tightly its electrons are held to the positively
charged nucleus and the more difficult they are to remove.
Electronegativity:
Electronegativity measures the ability of an atom to attract to itself the
electron pair forming a covalent bond.
The greater the electronegativity of an atom, the greater its affinity for
electrons.
Among the main-group elements, electronegativity
increasing
moving from left to right in the periodic table. Ordinary, it decreases
moving down a group.
SUMMARY:
Ionization energy & Electronegativity
Atomic radius &
Li Be
Ionization Energy
Na
and
K
Electronegativity
Rb
Cs
B
Ionic Radius
C N
Atomic Radius
and
Ionic Radius
Hybridization of atomic orbitals:
For example, Beryllium has atomic number of 4 and the electronic
configuration of them is 1s2 2s2
Fluorine has atomic number of 9 and the electronic configuration of them is
1s2 2s2 2p6
The formation of the BeF2 molecule can be explained by assuming that, as
two fluorine atoms approach, the atomic orbitals of the beryllium atom
undergo a significant change. Specifically, the 2s orbital is mixed or hybridized
with a 2p orbital to form two sp hybrid orbitals.
one s atomic orbital + one p atomic orbital = two sp hybrid orbital
Notice that the number of hybrid orbitals formed is equal to the number of
atomic orbitals mixed. Also, the energies of the hybrid orbitals are
intermediate between those of the atomic orbitals from which they are
derived.
Hybrid orbitals and their geometries:
Number of
Atomic orbitals
electron pairs
Hybrid
orientation
Examples
orbitals
2
S,p
Sp
Linear
BeF2 CO2
3
s, two p
sp2
Triangular planar
BF3 SO3
4
s, three p
sp3
tetrahedron
CH4 NH3 H2O
5
s, three p, d
sp3d
Triangular bipyramid
PCl5 SF4
6
s, three p, two d
sp3d2
octahedron
SF6 XeF4