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CHAPTER 1
ATOMIC STRUCTURE
1. INTRODUCTION
It was in the fifteenth-century that the Greek philosopher Democritus expressed
the belief that all matter is composed of small, indivisible particles that is called atomos
Although many of his contemporaries did not accept his idea, his suggestion persisted
throughout the centuries. Evidence from early experiments supported the notion of
“atomism”, but it was not until less that an English chemist, mathematician and
philosopher, John Dalton formulated a precise definition of the indivisible building
blocks of matter that we call atoms.
2. THE STRUCTURE OF THE ATOM
2.1 Dalton’s Atomic Theory
Dalton’s atomic theory is composed of the following hypotheses about the nature
of matter:
1. All matter is composed of atoms which are extremely small and indivisible particles.
The atom is the basic unit of an element that ----- enter into chemical combination.
2. All atoms-of-a given element are identical with the same size, mass, and chemical
properties. However, atoms of different elements differ in the said properties.
Example: All atoms of oxygen are identical. They have the same size, mass and chemical
properties. Atoms of nitrogen are different from of oxygen. The mass, size
and chemical properties of an oxygen atom are different from those of
nitrogen atom.
3. Atoms of different elements combine with each other in a definite ratio of small
whole numbers to form the same compound. (Law of Definite Proportion). A
compound always has the same composition regardless of the quantity or source.
Example: Whenever hydrogen atoms combine with oxygen atoms to form water, they
always combine in the ratio of two hydrogen atoms to one oxygen atom.
Hence, the formula for water is always written as H2O.
4. During a chemical reaction, atoms are either combined, united or arranged; some are
created or destroyed (Law of Conservation of Mass). This is the basic of a balanced
equation.
Example: In the combustion of carbon to carbon dioxide,
C(s)  O2 ( g )  CO2 ( g )
one carbon atom and two carbon atoms combine to form the product. There is
no increase or decrease in the number of carbon atoms or oxygen atoms
during the reaction. The number of carbon and oxygen atoms present before
reaction is the same as that present after the reaction has taken place.
2.2 The Subatomic Particles
Although an atom is indivisible, atoms are actually made up of even smaller
particles subatomic-particles. There are three subatomic particles – electrons, protons and
neutron.
Table 1.1. Mass and charge of subatomic particles.
Mass
Charge
Particles
Grams
Atomic mass unit
Cculomb
Charge Unit
-28
-24
Electron
9.1095 x 10
0.000549
-1.6022 x 10
-1
Proton
1.67252 x 10-24
1.007260
+1.6022 x 10-19
+1
T
Neutron
1.67495 x 10-24
1.008670
0
0
a
Table 1.1 summarizes the mass and charge of the three principal subatomic particles that
are important in chemistry – the electron, the proton and the neutron.
2.3 Atomic Number, mass Number and Isotopes
All atoms can be identified by the number of protons and neutrons.
The atomic number (Z) is the number of protons in the nucleus of each atom of an
element. This is a formed number and is a characteristic of the element.
The mass number (A) is the total number of protons and neutrons present in the
nucleus of an atom of an element. This number may vary from atom to atom and even
among atoms of the same element.
Therefore,
Mass number = number of protons + number of neutrons
= atomic number + number of neutrons
In a neutral atom, the number of protons equals the number of electrons. Hence,
atomic number equals the number of electrons.
A change in the distribution of electrons occurs during chemical reactions. In this
case of charged atoms (ions), the number of protons is uncharged because chemical
reactions involved orbital electrons only. The number of electrons can be either lower
than the number of protons because the atom lost electrons (the ion is positively charged)
or higher than the number of protons (when atom is negatively charges) then the number
of protons.
Charge = algebraic sum of positive charges and negative charges
An atom may therefore be represented this way: ZA Symbol
Example: Let us consider the 1123 Na atom.
Since the atomic number of sodium (Na0 is 11, there are 11 protons in the Na
atom. The mass number of 23 means that there are 23 nucleons (protons and
neutrons taken collectively). But since there are 11 protons, there must be (23-11)
12 neutrons. The atom has 11 electrons because it is neutral.
Example: Take the case of the 1224 Mg2+ ion.
The atom has 12 protons (Z = 12), 12 neutrons (n=24-12) and only 10 electrons
(+2 = +12 – 10). The charge of (+2) is due to the loss of 2 electrons.
Test Yourself:
Supply the missing data to complete the following table:
Symbol
A
Z B
32 A216 S
107
Ag+
47
27
n+
Z Sym
A
n17 Sym
Z
A
no. of protons
5
No. of neutrons
6
no. of electrons
18
The atomic number, Z, of an atom is characteristics of the element. Such cannot
be said of the mass number, A, because atoms of the same element (having the same
number of protons) can have different number of neutrons. Atoms of the same element
(same Z and therefore equal number of protons) which have different number of neutrons
(and therefore different A) are called isotopes. The properties of the isotopes of Ne (z =
10) are given in Table 1.2
Table 1.2. Properties of the isotopes of Neon.
Isotopic
Symbol
1020 Ne
1021 Ne1022 Ne+
Atomic
Number
10
10
10
Nuclear
Charge
+10
+10
+10
Mass
Number
20
21
22
No. of
protons
10
10
10
No. of
neutrons
10
11
12
Isotopic
Mass(amu)
19.99
20.99
21.91
The listed atomic masses do not correspond to the isotopic mass of any one
isotopic because they are determined from the average of the isotopic masses of naturally
occurring isotopic. The average is determined by taking into account the relative
abundance of the different isotopes.
Example: the isotopes of Si have the following isotopic masses and relative abundance:
Isotope
1428 Si
14298 Si
1430 Si
Isotopic Mass( amu)
27.9866
28.9866
29.9832
Isotopic Abundance (%)
92.28
4.67
3.05
The average atomic mass is determined as follows:
Average atomic mass = 0.9228(27.9866) + 0.0467(28.9866) + 0.0305(29.9832)
= 28.3 amu
Test Yourself:
Determine the average atomic mass of Mg from the following data:
Isotope
1224 Mg
1225 Mg
1220 Mg
Isotopic Mass
23.9924
24.9938
23.9895
Relative Abundance (%)
78.6
10.1
11.3
2.4 Development of the Atomic Model
The following works contributed significantly to the development of the Atomic Model:
1. Dalton’s “Billiard Ball” Model: Atoms are hard indestructible sphere.
2. Thomson’s “Raising Bread Model”: An atom is a sphere of positive particles – protons
(bread) in which are imbedded the negative particles – electrons (raisin).
3. Rutherford’s Nuclear Model: Most of the mass and positive charges of an atom are
concentrated in a very small region in the center of the atom called the nucleus. The
electrons move in the outer space which comprises the rest of the volume of the atom.
4. Bohr Model of the Hydrogen Atom: The electron moves about the nucleus in circular
orbits or energy states which possess the following properties:
a. Only certain orbits (and therefore certain energies) are allowed.
b. Energies associated with motion of electrons in the said orbits have fixed values
(quantived).
c. Each orbit has an assigned principal, quantum, number, n whose values are
restricted to whole number: n = 1,2,3,4, etc. the energy of an electron in a higher
energy level (n>1) can be determined by using equation:
E n   RH (1 / m 2 )
Where: RH, Rydberg constant, has value of 2.1 x 10-18J.
According to the equation, as the value of m decreases, the absolute value of En
increases but becomes more negative, with the most negative value at m = 1 , the most
stable orbit.
Example: Calculate the energy of an electron in the third main energy level.
Using the above formula we get the energy of the electron for n = 3:
En = -2.1 x 10 -18 (1/32) = -2.1 x 10-18 (1/9) = 2.33 x 10-19 J
The model further allows the electron to remain in orbit indefinitely without emission of
energy. Note that the second property of the energy states in that their energies are quantized.
In the presence of radiant energy, the electron may absorb energy and “jump” to a higher
orbit. On the other hand, an electron in a higher orbit may omit energy and “drop” to a lower
orbit. The energy that the electron absorbs or omits is in planck’s, as quantum. The energy of a
packed or quantum is equal to the difference in energy of the two orbits between which the
transition (jump or drop) -------.
The difference in energy may be determined by the equation:
E  Ef  Ei

 
  RH (1 / nf 2 )   RH (1 / ni )
2

  RH (1 / nf 2 )  (1 / nf 2 )
The wavelenght of the photon emitted by the electron can then be determined by the
formula:
E  hc / 
Where: h (Plank’s constant) = 6.63 x 10-34 s
c (speed of light) = 3.00 x 108 M/S
 wavelenght
Example: Calculate the energy associated with the process when an electron “drop” from the ni =
5 state to the nf = 2 state in the H atom. Determine also the wavelength of the emitted
photon.
Using the formula for  we get
 = 2.1 x 10-18 ( 1/518 -1/22 )
= 2.1 x 10-18 ( 1/25 - 1/4 )
= - 4.41 x 10-19 J
 = hc / λ
-19
-34
x 10 J = (6.63 x 10 s x 3.00 x 108 m/s / λ
λ = 4.51 x 10-7 m
= 4.51 x 10-7 m x (1 x 109 nm/m)
= 451 nm
Test Yourself:
Calculate the wavelength of the photon emitted during a transition from the ni = 6 to the
nf – 3 state in the H atom.
5. Quantum Mechanical Model: The electron moves in a three dimensional space
around the nucleus.
3. QUANTUM NUMBERS
In quantum mechanics, three quantum numbers are used to describe the
distribution of electrons in atom – the first three quantum numbers – the principal
number, the angular momentum quantum number and the magnetic quantum number –
describe the energy of the electron while the fourth quantum number – the spts quantum
number describes the behavior of a specific electron and completes the description of
electrons in atoms.
3.1 The Principal Quantum Number (n)
The principal quantum number (n) is primarily used for determining the overall
energy of an atomic orbital. It corresponds to the main energy level or shell and can have
integral values 1, 2, 3, and so forth. It also relates to the average distance of the electron
from the nucleus in a particular orbital. The large n is, the greater is the average energy of
the shell or the average distance of an electron in the nucleus and therefore the larger is
the orbital.
3.2 The Angular Momentum Quantum Number (t)
The angular momentum quantum number, t, determines the angular momentum
and the shape of the orbital. It has a smaller effect on the energy and corresponds to the
sublevel or enbshell. The values of t depend on the value of the principal quantum
number n and the ranges from 0 up to (n-1). The t quantum number also have
corresponding letter designations of s(t=0), p(t=1), d(t=2) and f(t=3).
The number of sublevels (subshell) in any given main energy level (shell) is equal
to the values of n.
To specify a sublevel (subshell) within a given main energy level (shell), we write
the values of n for the main energy level followed by the letter designations of the
sublevel.
Example: The p sublevel of the second main energy level is designated 2p.
3.3 The Magnetic Quantum Number (mt)
The magnetic quantum number mt determine the orientation of the angular
momentum vector in a magnetic field and describes the orientation of the orbited in
space. For every sublevel, the value of – depends on the value of the angular momentum
quantum number t and has values which range from –t passing through zero up to –t. H
Hence for t = l (a p sublevel), the possible value of t are -1, 0 and +1.
The number values of – which is the number of orbitals in the sublevel is equal to
(2-t---). The total number of orbitals in an energy level is given by n2.
3.4 The Electron Spin Quantum Number (ms)
The electron spin quantum number, ms determines the orientation of the electron
magnetic moment in a magnetic field, either in the direction of the field (+1/2) or
opposed to it (-1/2). Therefore, there are only two positive values of the electron spin
quantum number, -1/2 or +1/2. This is given by Rexli’s exc---- principle which states that
no two electrons can have exactly the same ‘set” of quantum number. Hence, an orbital
can only accommodate two electrons of opposite spins.
The quantum numbers and their properties are summarized in Table 1.2
Table 1.2. Quantum numbers and their properties
Symbol
Name
Values
Role
n
Principal
1,2,3…..
Determine the major part of the
energy
t
Angular
momentum
0,1,2,3,…(n-1)
Describe “angular dependence and
contributed to the energy
mt
Magnetic
0, ±1, ±2,.. ±1
Describes orientation in space
ms
Spin
±1/2
Describes orientation of electron
spin in space
Test Yourself
Complete the following table:
n
t
mt
Letter
values
values
Designation
1
2
3
4
5
No. of
sublevels
No. of orbitals
In sublevel
In energy level
Test Yourself:
For each set of quantum numbers given below, indicate whether the combination is
possible or not. If the combination is not possible, state the reason why.
a.
b.
c.
d.
4
n = 4,
n = 5,
n = 6,
n = 2,
t = 3,
t = 1,
t = 0,
t = 2,
m2 = +3
m2 = +0
m2 = +1
m2 = 0
ATOMIC ORBITALS
In principles an electron can be found anywhere but most of the time it is quite close
to the nucleus. If electron density in an orbital is plotted against distance from the nucleus, it
will be observed tat electron density decreases as the distance from the nucleus increases.
Hence an orbital is referred as the region in space where there is highest probability of
finding an electron.
4.1 The s Orbital
All s orbitals are spherical in shape. They differ in size which increases as the
principal quantum number increases.
4.2 The p Orbital
The letter subcrippts in the designation of the p orbitals indicate the axes along which
the orbitals are oriented. These three p orbitals are identical in size, shape and energy
(degenerate). They differ into one anther only in their orientations and may be designated as
px,py and pz to indicate the axes of orientation of the orbitals. There is no simple relation
between the values of mt and x, y and z directions. Like s orbitals, p orbitals increase in size
as the principal quantum number increases.
4.3 The d Orbitals
Just like p orbitals, there is no direct relation between a given orientation and mt
value although the different orientations of the d orbitals corresponds to the different values
of mt. All d orbitals in the same sublevel are degenerate and may be designated as
d xy , d xz , d yz , d x 2 y 2 and d x 2 . The first four orbitals have four lobes with the nucleus by the
center, while the d z 2 orbitals has two lobes with a collar where the electron density is
highest.
4.4 f Orbitals
Just like p and d orbitals, there is also no direct relation between a given orientation
and an mt value in f orbitals. There are seven f orbitals with mt value ranging from -3 passing
through 0 up to +3.
5
THE ENERGIES OF ORBITALS
The energy of an electron in a many electron atom depends not only entire principal
quantum member but also on the angular momentum quantum number. Hence, the
distribution of electrons in an atom may be predicted by taking into account the n and t
quantum numbers (n + t rule). As the value of (n + t) increases, the energy of the electron
also increases.
Applying the rule, we can deduce that the energy of a 3p(n + t = 3 + 1 = 4) electron is
greater than that of a 3s (n + t = 3 + 0 = 3); 3d (n + t = 3 + 2 = 5) has higher energy than 4s(n
+ t = 4+0 = 4); 4f (n + t = 4+3 = 7) has higher energy than 6s (n + t = 6+0 = 6). In cases
wherin the (n + t) values are equal, the orbital in the higher energy level has the higher
energy. This is the reason why although 4s (n + t) = 4) and 3p (n + t =4) have the same (n + t)
value, the 3p sublevel in filled first before the 4s sublevel.
5.1 Electron Distribution
The following rules are useful in predicting the electron distribution in an atom:
1. n + t rule. As the (n + t) value increases, the energy of the electron increases.
2. Autbau (building up). The buildup of electron in atoms results from continually
increasing the quantum numbers, starting with the lowest n, t and mt values.
Applying rules 1 and 2, we arrive at the following diagram:
where the direction of the arrows indicate the order of filling up of sublevels.
3. Hand’s rule of maximum multiplicity. Electrons entering a sublevel containing more
than one orbital (degenerate) will be appeared cut with their epics in the same
direction so that electrons are distributed singly before they are painted. Hence, no
orbital can have a pair of electrons, if other orbitals in the same sublevel are empty.
To represent the electron distribution of C (Z = 6), we write 1s2 2s2 2p2. The
following orbital diagram can be represented
or
1s
2s
2p
2p
2p
1s
2s
2p
2p
2p
or
1s
2s
2p
2p
2p
because the three p orbitals are degenerate.
Example: Predict the distribution of electrons and the quantum numbers of all electrons
in oxygen (Z = 8).
Solution:
The electron distribution of gO is given by: 1s2 2s2 2p4 and the orbital diagram is
represented by
1s
2s
2p
2p
2p
The quantum numbers can therefore be assigned as follows:
1
2
2
2
2
1
0
0
1
1
1
mt
0
0
-1,-1
0
+1
-1/2
+1/2
-1/2
+1/2
-1/2
+1/2
-1/2
-1/2
ms
Test Yourself:
Determine the quantum numbers of the 27th electron of Co.
SUMMARY
There are three principle subatomic particles – the electron, the proton and
neutron. This proton is positively charged and has a mass number of one, the neutron is
neutral and has mass number of one; the electron is negatively charged and has a mass of
zero.
The differences in the number of subatomic particles are responsible for the
existence of isotopes, isobars and isotones.
Three quantum numbers are used to describe the distribution of electrons in an
atom: the principal quantum (n), the angular momentum quantum (t) and the magnetic
quantum number (mt). The orientation of the electron magnetic moment is determined by
the magnetic number (ms).
The following rules apply in predicting the distribution of electrons in an atom:
the (n + t) rule, the Autbau building up principle and Hund’s rule of maximum mutlipicit.
CHAPTER 2
THE PERIODIC TABLE
1. INTRODUCTION
Mendeleov’s and Meyer organized the elements in order of atomic weight then
identified facilities of elements with similar properties. By arranging the families in rows and
columns and considering similarities in chemical behavior as well as atomic weight.
Mendeleov’s was able to predict the properties of several elements which have not yet been
discovered. The discovery of additional elements not known in Mendeleov’s time and the
synthesis of heavy elements have led to the more complete periodic table in its modern form.
2. ATOMIC STRUCTURE AND THE PERIODIC TABLE
The position of an element in the periodic table is defined by periods and group or families.
A horizontal row of elements is called a period, and a vertical column in a group. The period
number is related to the highest value of the principal quantum numbers. The group number is
related to the differentiating electron.
Hence,
Period number = highest value of n
= number of occupied main energy levels
The group number, on the other hand, can be determined as follows:
Using the previous notation (U.S. convention), the family may be determined as
follows:
Family = A (if last electron is in s or p sublevel)
= B (if last electron is in d or f sublevel)
The group number can be determined by using the following formula.
Group number = number of valence electrons (if A family)
= number of [ ns + (n-1) d electrons (B family)
Where n is the period number
Example: Consider AI (Z=13) with electronic configuration of: 1s2 2s2 2p6 3s2 3p1
AI belongs to period 3 because there are three occupied energy levels. Since the last
electron is in the s sublevel it belongs to the A family. The group number of an element in an the
A family is equal to the number of valence electrons. AI has 2 electrons in 3s and one electron in
3p giving a total of three. AI is therefore found in period 3, Group IIIA.
Example: Consider V(Z-23) with electronic configuration of: 1s2 2s3 2p6 3s2 3p6 4s2 3d3 V
is in period 4 became there are four occupied energy levels. Since the last electron is in the 3d
sublevel, it belongs to the B family. For an element of the B family, the Group Number may be
determined as follows:
Period number = 4, therefore n = 4
Using the formula for group number of a B family element.
Group number = number of [ns +(n+1) d electrons
We get
Group number = 2 electrons in 4s+3 electrons in 3d = V
V is in period 4, Group VB.
3.1 IUPAC Notation: The s- and p- Blocks.
Groups 1 and 2 are referred to as the s-blocks because the last electron is in the s orbitals
of the outermost energy level of the atom. Since the s orbital can only accommodate 2 electrons,
there are only 2 groups in the s-blocks group 1 and 2.
Group 13 to 18 are called the p-blocks because the last electrons goes to a p sublevel.
There are 6 groups in this block because the p orbital can accommodate a maximum of 6
electrons.
For s- and p-blocks elements, the following rule is used to determine group number:
s-blocks
Group number = number of valence electrons
p-blocks
Group number= 10+ number of valence electrons
Note that:
Unit, digit of group number = number of valence electrons
In the IUPAC notation, electron distribution and group number are related as follows:
1
2
13
14
15
16
17
18
Group number
Ns1
Ns2
Ns2np1 Ns2np2 Ns2np2 Ns2np4 Ns2np5 Ns3np6
General electrons
configuration
1
2
3
4
5
6
7
8
Number of valence
electrons
Example: Using the IUPAC notation, AI (in the previous example) with 3 valence electrons
belongs to period 3, Group 13.
3.2 The d- Transition Elements
The d-blocks of elements (those which correspond to the filling of the d orbital) belongs
to Group 3x12. There are ten groups in this block because there are ten-electrons in a d sublevel.
For these groups of elements the group number is determined by the formula.
Group number = number of valence electrons
Where: number of valence electrons = number of ns+(n-1)d electrons with
a=period time
Example: To determine group number of V ( electron distribution of [Ar]4s2 3d3) using formula:
Group Number = number of 4s electrons + number of 3d electrons
= 2+3 = 5
3.3 The f-Transition Elements
Below the s-, p- and d-blocks are two rows of elements called the f-transition elements.
The first row which is found is period is called the leuthunide varies-or- lenthankles or are earth
elements. The second row, in period 2, is called the methaides. Each row consists of 14-elements
because there are 14 electrons involved in the filling of the seven 5f or 6f orbital. The valence
electrons of these elements are in the
orbitals where only the the number of (n-1)d
and (n-2)f electons vary. Since these elements vary only in the number of electrons in their inner
main energy levels, their chemical properties are very similar.
Supply the missing information without looking or your periodic table.
Elements
Valence of
Period Number
Group Number
Electrons
US Notation
IUPAC Notation
11Na
20Ca
29Cu
52Tc
50Sn
25Mn
4. PERIODIC PROPERTIES OF ELEMENTS
The electrons distribution of the elements show a periodic with increasing, atomic
number. As a result, elements also exhibit periodic variation in their physical and chemical
behavior.
4.1 Atomic Radius
Atomic radius is one-half the distance between two nuclei in two adjacent metal atoms or
for elements that exist as simple diatomic molecules, it is one-halt-the distance between the
nuclei of the two atoms in a particular molecule.
The general trend across a period is a decrease in atomic size as nuclear charge increase.
The atomic radius is determined to a large extent by the strength-of-attention between the
outer shell electrons and the nucleus. The larger the effective nuclear charge, the more strongly
held the outer electrons are and the smaller the atomic radius is.
Consider the second-period elements. For all elements is this period, as the atomic
number increase, the number of electrons in the inner shell remains constant while the nuclear
charge increase. The electrons that are added to counterbalance the increasing nuclear charge are
ineffective in shielding one another because they are in the same energy level. Hence, the
effective nuclear charge increases steadily while the principal quantum number remains constant.
Example: Take the case of lithium (Z-3 and electronic configuration is 1s2 2s1). The
shielding effect of the two 1s electron is to cancel two positive charges in the nucleus, leaving
one positive charge uncancelled. The 2s electron only feels the attraction due to one proton (that
whose positive charge was not cancelled by shielding electrons) in the nucleus, or the effective
nuclear charge is +1.
If we consider beryllium (Z=4, 1s2 2s2), each of the 2s electrons cancel two of the four positive
charges in the nucleus and the effective nuclear charge is +2 because the 2s electrons do not
shield each other. Therefore, be atom is smaller than Li stone.
On the other hand, as Z increase within a group, the atomic radius increase. Remember
that, as the principal quantum number n increase orbitai size increase. Hence, atomic radius
increases as n increase.
4.2 Ionizations Energy
Ionization energy, also known as ionization potential, is the minimum energy required to
remove electron from an isolated gaseous atom or ion in its ground state.
An+ (g)  A(n+1)+ (g) +e
Ionization energy =  
n+
Or
A (g) +    A(g)+e
For an atom with several electrons, the amount of energy required to remove the first
electrons from the atom in its ground state is called first ionization energy.
A(g) +   1 A1+(g) e
The second ionization energy (   2) and the third ionization energy (   3) are shown
by the following equations:
A1+(g) +   2 A2+(g)+e
A2+(g) +   3 A3+(g)+e
After an electron is removed from a neutral atom, the repulsion-among the remaining
electrons-therefore but since nuclear charge remain constant, more energy is used is remove
another electron from the positively charged ion. Therefore, ionization-energies-always increase
as the charge-increase.
  1<   2<   3
The general trend across a period is an increase in ionization energy the nuclear charge
increase with the…………………….
Example: In the second period, the noble gas Ne has the highest ionization, while Li, the
first elements in the period, has the lowest.
Breaks in the trend, that is, a decrease in ionization energy appear at the following points.
1) Upon the addition of the first na electrons. The new electrons is in an orbital that has
most of its electron density further away from the nucleus than the second as
electrons.
Example: AI([Ne] 3s2 3p1) would expected to exhibit a higher ionization energy than
12Mg ([Ne] 3s2) because AI comes after (reading from left to right)Mg in the periodic
table. However, the ionization energy of AI (   1=577.6 k3/mol) is observed to be lower
than of Mg(…)
2) Upon the addition of the fourth … electron. The new electron shares an orbital with
mother electron and the pairing energy or repulsion between two electrons in the
same region of space, reduces the ionization energy.
Example: Based on the observed trend … shol exhibit a higher ionization energy than…
However the actual values, shows that S (….) has a lower ionization energy (…) than P(…)
Similarly patterns appear in lower periods, with only small changes in ionization energy
for the transition elements, and generally lower value for heavier atoms in the same family
because of increased shielding by inner electrons and increased distance between the… and the
outer electrons.
Much larger decreases in ionization energy occur at the start of each new period, because
the change to the next major quantum number requires a higher energy. The maxima at the noble
gases decrease with increasing Z, because the outer electrons are farther from the… in the
heavier elements.
4.3 Electron Affinity
Electron affinity is the energy that is released or absorbed when an electron is added to a
gaseous neutral atom in its ground state.
A(g)+e  A(g)
The more greater the electron affinity, the greater the tendency of the atom to accept an electron
The overall trend is that the tendency to accept electrons…(that is, the electron affinity
values increases) as the atomic number ………… a period. This means that the electron affinities
of metals are generally lower(or more negative) than those of nonmetals.
The following electron affiliations are observed with elements in the second period.
Elements Na
Electrons 52.9
affinity
(kl/mol)
Mg
-39
AI
42.6
Si
133.6
P
72.0
S
200.4
CI
349.0
Ar
-97
Among the elements within a given period, the halogen kits the greatest electron affinity
because it is the most nonmetallic. However, the show little variation a group.
SUMMARY
The position of an elements in the periodic table in defined in terms of the period and the
group.
Elements display periodic variation in their physical and chemical behavior.
Atomic radius is observed to decrease with atomic number within a period and to
increase with atomic number within a group.
Ionization energy increase with increase in atomic number within a period and decrease
with increase in atomic number within a group.
Electron affinity is observed to become more negative as the tendency to attract electrons
increase.