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The elements on the periodic table exhibit gradual variation in the following properties
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Atomization
2
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Ionization
4
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Electronegativity is the same as the enthalpy of atomization
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They would have the same chemical and physical properties
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That is to say, the atomic number is the same as the proton number which is equal to the
number of electrons around the nucleus in a neutral atom
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Atomic
orbital is the maximum volume of space around the nucleus in which there is a higher probability of
locating an electron
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An electron in shell “K” has
different energy from that of an electron in the “Q “shell
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The way electrons are distributed in an atom is known as the electron configuration
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The electrons in the
outermost Shell are called the valence electrons and the shell is a valence shell
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The core electrons are also known as the inner electrons and
they are located between the nucleus and the valence electrons they do not participate in chemical
reactions but do influence the behaviour of valence electrons
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Meaning that, the properties of
the various elements on the periodic table are a result of the difference in atomic numbers
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The periodic table is simply a table showing all known elements chronologically in order of
increasing atomic number such that the outer electron configuration and properties recur at intervals
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Numbers are
assigned to a group depending on the number of electrons in the outermost shell of the atom
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It has a number equal to the
number of shells in the atom containing elections

The Alkali Metals- Group 1 elements
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They are
univalent that they have one electron in their outermost shells and they form an ion by donating the
one electron in their outermost shells
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• They are very reactive and normally found in nature combined with other elements
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They are therefore
stored in inert liquids an example of such liquids is paraffin oil
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• The malleable that is they can easily be beaten into shapes and ductile that is they can be drawn
into thin wires
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The Alkaline Earth Metals: Group 2
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•
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•
•
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Ionize by giving off two electrons
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They are less reactive than alkali earth metals
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They are good conductors of heat and electricity
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•
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•
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They are Fluorine, Chlorine Bromine Iodine, and Astatine
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They react to gain one electron easily to form anions with an oxidation number of -1
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They do not conduct electricity because they are strong
oxidizing agents
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They include Helium, Neon, Argon Krypton Xenon and
Radon
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• They are generally chemically inactive
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Helium is used to fill meteorological balloons
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They are placed between groups 2 and 3 in period 4 on the periodic table
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These elements are Scandium, Titanium, Vanadium, Chromium,
Manganese, Iron, Cobalt, Nickel, Copper and Zinc
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Characteristics of transition Metals
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•
They show paramagnetic properties thus, they have their “d” orbitals partially filled and di- magnetic
means the “d” orbital is filled
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Hydrogen can be placed in both groups 1 and 7 because of its properties which resemble those of
groups 1 and 7
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Metalloids
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They are called semimetals and they are rather a mixture of
both metals and nonmetals
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• They are not good conductors of electricity but when slightly made impure they become good
conductors of electricity
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Physical Properties of metalloids
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Metallic character
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Metals
that lose electrons easily are said to be strongly metallic therefore exhibit greater metallic character
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Properties of metals
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Nonmetals
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Except for noble gas, non-metals
gain or share electrons to achieve an inert gas structure
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They are insulators that are they do not conduct heat and electricity
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Atomization energy
Atomic radius
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The size of an atom depends on:
• The nuclear charge (proton number)
• The number of the core-shell where core shells are the inner shells of an atom that comes
before the valence shell
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Two atoms having the same number of electrons, the one with a greater proton in the nucleus
will attract or pull the electrons in the outermost Shell more closely to itself which leads to a decrease in
size
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The attraction by the positive protons also called nuclear charge for the electrons in the
outermost shells is usually reduced due to the presence of core electrons in the core shells which cause
the shielding effect
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Therefore: an effective nuclear charge is a net positive charge which is felt by electrons in the
outermost Shell through attraction by the positive protons that are,
EFFECTIVE NUCLEAR CHARGE = (NUCLEAR CHARGE) – (SHIELDING EFFECT)
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This is because across the
period the nuclear charge increase likewise the number of electrons but the core shells do not increase
because the electrons fill the outermost shell which is not fully filled as a result, the nuclear charge can
pull the outermost shell more closely to a self reducing the atomic size of atoms of different elements
across the period
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The atomic radius of inert gas is higher than the elements across the period
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Example:
Arrange the following in order of increasing atomic size K, Li and Na
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Answer:
Li – Na – K
Increasing atomic size from left to right
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Therefore, since down the group, the various atoms of different elements follow
the pattern li Na and k down the group increase takes place the same way
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Reason:
With the loss of one or more electrons, it means there are fewer electrons left in the iron formed to be
pulled by the same number of protons as in the neutral atom
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Note: when a greater number of electrons are lost from an atom the remaining
electrons are attracted more strongly or pulled in nearer to the positive nucleus
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Anions: A negative ion or anion has a larger radius than a neutral atom
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The outer electrons become less tightly bound to the nucleus which leads to the increase in atomic size
or radius of the anion
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An effective nuclear charge cannot Pull the
electrons in the anion as tightly as in the neutral atom
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Note:
Across the period and Ion decreases due to increasing effective nuclear charge
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Isoelectronic Series
Is this series of ions and atoms with the same number of electrons
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Ionization energy
It is the minimum energy required to remove an electron from the outermost shell of a free or
isolated gaseous atom to form a gaseous singly charged cation
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The second ionization energy of an element is the energy required to remove one mole of electrons
from one mole of gaseous singly charged cation
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Mg2+ -> Mg2+ + e(electron) 3rd Ionization energy
Factors influencing ionization energy
Atomic radius or the distance of the outermost electron from the nucleus
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•
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How stable electron configuration is
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1
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Reason:
Down the same group atomic radius increases
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Therefore less
energy is required to remove outer electrons down the group period this is because the outermost
electrons are not tightly held or pulled by the nuclear charge which results in an increase in size down
the group
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Ionization energy increases across the period from left to right
Reasons:
Across the period positive nuclear charge increases without the addition of any extra electron
shells even though the electrons do increase, they still fill the outermost shell which is incomplete
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As a result, atomic radius decreases as the electrons are held more firmly by the increasing
nuclear charge
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The first ionization energies do not increase smoothly across a period
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Notes: higher energies are required to remove electrons from their filled outermost shell due to stable
electronic structure
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Reasons:
With the loss of electrons, there will be fewer electrons left in the outermost Shell
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Since the outermost
shell electrons are strongly held more energy will be required to remove them from the ion
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Greater energy will then be needed to remove outer shell electrons from the more stable
ion
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A
First electron affinity is the energy lost or gained when one mole of atoms in a free gaseous State gain
one mole of an electron into the outermost shells to form one mole of gaseous singly charged anion(s)
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Factors influencing Electron Affinity
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✓ Down the group electron affinity decreases
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Less energy is therefore released when an electron is added

Notes:
Down group 7 elements, Fluorine has unexpectedly lower electron affinity
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The electron has to be forced in and this reduces the overall energy
released
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Reason:
Atomic size decreases together with a corresponding increase in effective nuclear charge
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While the first electron affinity for most electronegative elements is negative, why are the second
electron affinity values for the singly charged anions happen to be positive?
Reasons:
The negative charge due to the second electron affinity creates a very strong repulsive field for the
incoming negative election
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A large amount of energy has to be absorbed to do that and hence a positive electron
affinity value for the second electron affinity
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Electronegativity, therefore, increases across the
period from left to right as effective nuclear charge increases with decreasing size
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In a bonded molecule the more
electronegative atom acquires a negative charge whiles the less electronegative atom obtains a positive
charge
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Noble gases do not have electronegativity due to the stable nature of the
electron configuration
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It is the energy absorbed when 1 mole of free gaseous atoms are formed from the elements at
standard conditions (298kelvin, and 100kpa pressure)
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Sublimation occurs when a solid absorbs heat and changes directly into the
gaseous state without first changing into liquid
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Covalent substances undergo dissociation
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Dissociation energy is twice the atomization energy for diatomic
molecules
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(gaseous) -> (gaseous)

Cl2 -> 2Cl2 ∆H diss = +241kJmol-1
(gaseous)->(gaseous)
When a bond is strong its atomization energy is very high
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