The Periodic Table

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Chapter 4 The Periodic Table
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Chapter 4 The Periodic Table
Section 1 - How Are Elements Organized?
Pure elements at room temperature and atmospheric
pressure can be solids, liquids, or gases. Some
elements are colorless. Others are colored.
Despite the differences between elements, groups
of elements share certain properties.
Similarly, the elements fluorine, chlorine,
bromine, and iodine can combine with sodium in a
11 ratio to form NaF, NaCl, NaBr, and NaI. These
compounds are also white solids that can dissolve
in water to form solutions that conduct
electricity. These examples show that even
though each element is different, groups of them
have much in common.
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Chapter 4 The Periodic Table
Section 1 - How Are Elements Organized?
In 1865, the English chemist John Newlands
arranged the known elements according to their
properties and in order of increasing atomic
mass. He placed the elements in a table.
In 1869, the Russian chemist Dmitri Mendeleev
used Newlandss observation and other information
to produce the first orderly arrangement,
or periodic table, of all 63 elements known at
the time. Mendeleev arranged the elements in
order of increasing atomic mass. Mendeleev
started a new row each time he noticed that the
chemical properties of the elements repeated.
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Chapter 4 The Periodic Table
Section 1 - How Are Elements Organized?
He placed elements in the new row directly below
elements of similar chemical properties in the
preceding row. He arrived at the pattern shown
below.
Two interesting observations can be made about
Mendeleevs table. First, Mendeleevs table
contains gaps that elements with particular
properties should fill. He predicted the
properties of the missing elements.
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Chapter 4 The Periodic Table
Section 1 - How Are Elements Organized?
Second, the elements do not always fit neatly
in order of atomic mass. For example, Mendeleev
had to switch the order of tellurium, Te,
and iodine, I, to keep similar elements in the
same column. At first, he thought that their
atomic masses were wrong. However, careful
research by others showed that they were correct.
Mendeleev could not explain why his order was not
always the same.
About 40 years after Mendeleev published his
periodic table, an English chemist named Henry
Moseley found a different physical basis for
the arrangement of elements. Moseley realized
that the periodic table should be arranged by
atomic number, not atomic mass. When the elements
were arranged by increasing atomic number, the
discrepancies in Mendeleevs table disappeared.
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Chapter 4 The Periodic Table
Section 1 - How Are Elements Organized?
The periodic law states that the repeating
physical and chemical properties of elements
change periodically with their atomic number
To understand why elements with similar
properties appear at regular intervals in the
periodic table, you need to examine the electron
configurations of the elements.
Elements in each column of the table have the
same number of electrons in their outer energy
level. These electrons are called valence
electrons.
A vertical column on the periodic table is called
a group or family.
A horizontal row on the periodic table is called
a period.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
Elements in groups 1, 2, and 1318 are known as
the main group elements.
Four groups within the main-group elements have
special names. These groups are the alkali metals
(Group 1), the alkaline-earth metals (Group 2),
the halogens (Group 17), and the noble gases
(Group 18).
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
The Alkali Metals Make Up Group 1
Alkali metals are so named because they are
metals that react with water to make alkaline
solutions.
All have a single valence electron
Alkali metals are usually stored in oil to keep
them from reacting with the oxygen and water in
the air. Because of their high reactivity,
Alkali metals are never found in nature as pure
elements but are found combined with other
elements as compounds.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
The Alkaline-Earth Metals Make Up Group 2
Like the alkali metals, the alkaline-earth metals
are highly reactive, so they are usually found
as compounds rather than as pure elements
All have 2 valence electrons
The alkaline-earth metals are not as reactive as
alkali metals, however they are harder and have
higher melting points.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
The Halogens make up Group 17
Halogens have 7 valence electrons
The halogens are the most reactive group of
nonmetal elements
The halogens have a wide range of physical
properties. Fluorine and chlorine are gases at
room temperature, but bromine, is a liquid, and
iodine and astatine are solids.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
The Noble Gases make up Group 18
The noble gases were once called inert gases
because they were thought to be completely
unreactive.
The low reactivity of noble gases leads to some
special uses. Helium, a noble gas, is used to
fill blimps because it has a low density and is
not flammable.
Hydrogen Is in a Class by Itself
Hydrogen is the most common element in the
universe. It is estimated that about three out of
every four atoms in the universe are hydrogen.
Because it consists of just one proton and one
electron, hydrogen behaves unlike any other
element.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
The Transition metals make up Groups 3 through 12
Unlike the main-group elements, the
transition metals in each group do not have
identical outer electron configurations.
Generally, the transition metals are less
reactive than the alkali metals and the
alkaline-earth metals are.
Lanthanides and Actinides
Part of the last two periods of transition
metals are placed toward the bottom of the
periodic table to keep the table conveniently
narrow,
The elements in the first of these rows are
called the Lanthanides because their atomic
numbers follow the element lanthanum.
Likewise, elements in the row below the
lanthanides are called Actinides because they
follow actinium.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
Summary of properties of selected groups
Group 1 All have 1 valence electron, and will
easily lose this electron to form 1 ions
Group 2 All have 2 valence electrons, and will
easily lose both electrons to form 2 ions
Group 13 All have 3 valence electrons, and will
easily lose all of them to Form 3 ions
Group 16 all have 6 valence electrons and will
easily gain 2 electron to Form 2- ions
Group 17 all have 7 valence electrons and will
easily gain 1 electron to Form 1- ions
Group 18 have a full outer shell with 8
electrons and will not react with other atom
easily.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
Transition elements are able to have multiple
ion charges. All Possible charges are positive
due to lose of electrons. Many will form a
colored aqueous solution.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
Most elements of the periodic table are called
Metals
All metals are excellent conductors of
electricity.
Metals are excellent conductors of heat.
Metals are ductile and malleable.
Ductile means that the metal can be squeezed out
into a wire.
Malleable means that the metal can be hammered
or rolled into sheets.
Metals tend to lose electrons to complete the
octet rule.
All metals are solid except mercury, which is a
liquid.
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table
Most elements of the periodic table near groups
14 -18 are called Nonmetals
All nonmetals are poor conductors of electricity.
Metals are poor conductors of heat.
Solid nonmetals are brittle and break easily
Metals tend to gain electrons to complete the
octet rule.
Nonmetals are found in all three states of matter
with bromine being the only liquid.
Metalloids are elements that share properties of
both metals and nonmetals
Metalloids are elements B, Si, As, Te, At, Ge, Sb
and Po
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table Review
Questions
At STP which element has a definite shape and
volume? 1. Ag 2. Hg 3. Ne 4. Xe
Which element is a member of the halogen
family? 1. K 2. B 3. I 4. S
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table Review
Questions
The metalloids that are included in Group 15 are
antimony (Sb) and 1. N 2. P 3. As 4. Bi
A characteristic of most nonmetallic solids is
that they are 1. brittle 2. ductile 3.
malleable 4. conductors of electricity
More than two-thirds of the elements of the
Periodic Table are classified as 1. metalloids
2. metals 3. nonmetals 4. noble gases
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table Review
Questions
To which group do the alkaline earth metals
belong? 1. 1 2. 2 3. 11 4. 12
Which three elements have the most similar
chemical properties? 1. Ar, Kr, Br 2. K, Rb, Cs
3. B, C, N 4. O, N, Si
When metals form ions, they tend to do so by 1.
losing electrons and forming positive ions 2.
losing electrons and forming negative ions 3.
gaining electrons and forming positive ions 4.
gaining electrons and forming negative ions
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table Review
Questions
The chemical properties of the elements are
periodic functions of their atomic 1. masses 2.
weights 3. numbers 4. radii
Which halogen is a solid at STP? 1. fluorine 2.
chlorine 3. bromine 4. iodine
Which element is in Group 2 and Period 7 of the
Periodic Table? 1. magnesium 2. manganese 3.
radium 4. radon
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table Review
Questions
An atom in the ground state contains 8 valence
electrons. This atom is classified as a 1. metal
2. semimetal 3. noble gas 4. halogen
Which category is composed of elements that have
both positive and negative charges? 1. the
alkali metals 2. the transition metals 3. the
halogens 4. the alkaline earths
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table Review
Questions
Which element is so active chemically that it
occurs naturally only in compounds? 1. potassium
2. silver 3. copper 4. sulfur
The elements of the Periodic Table are arranged
in horizontal rows according to each successive
element's greater 1. atomic mass 2. atomic
radius 3. number of protons 4. number of
neutrons
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Chapter 4 The Periodic Table
Section 2 - Tour of the Periodic Table Review
Questions
In which list are the elements arranged in order
of increasing atomic mass? 1. Cl, K, Ar 2. Fe,
Co, Ni 3. Te, I, Xe 4. Ne, F, Na
Which Group of the Periodic Table contains atoms
with a stable outer electron configuration? 1. 1
2. 8 3. 16 4. 18
The element in Period 4 and Group 1 of the
Periodic Table would be classified as a 1. metal
2. metalloid 3. nonmetal 4. noble gas
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
The arrangement of the periodic table reveals
trends in the properties of the elements. A trend
is a predictable change in a particular direction.
Metallic Trend
All of the elements of the periodic table
increase in metallic characteristics as they get
closer to francium.
This will make group 1 the most metallic group
and period 7 the most metallic period.
Nonmetallic Trend
All of the elements of the periodic table
increase in nonmetallic characteristics as they
get closer to flourine.
This will make group 17 the most nonmetallic
group and period 2 the most nonmetallic period.
Note The noble gases are not considered part of
the metallic or nonmetallic trends due to there
resistance to chemical reactions.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Ionization Energy
The energy that is supplied to remove an electron
is the ionization energy of the atom.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Ionization Energy Decreases as You Move Down a
Group
Each element has more occupied energy levels
than the one above it has. Therefore, the
outermost electrons are farthest from the nucleus
in elements near the bottom of a group.
Similarly, as you move down a group, each
successive element contains more electrons in the
energy levels between the nucleus and
the outermost electrons. These inner electrons
shield the outermost electrons from the full
attractive force of the nucleus. This electron
shielding causes the outermost electrons to be
held less tightly to the nucleus.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Ionization Energy Increases as You Move Across a
Period
From one element to the next in a period, the
number of protons and the number of electrons
increase by one each. A higher nuclear charge
more strongly attracts the outer electrons in
the same energy level, but the electron-shielding
effect from inner-level electrons remains the
same. Thus, more energy is required to remove an
electron because the attractive force on them is
higher.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Atomic Radius
The exact size of an atom is hard to determine.
An atoms size depends on the volume occupied by
the electrons around the nucleus, and the
electrons do not move in well-defined paths.
Rather, the volume the electrons occupy is
thought of as an electron cloud, with no
clear-cut edge. In addition, the physical and
chemical state of an atom can change the size of
an electron cloud.
The diagram below shows one way to measure the
size of an atom. This Method involves
calculating the bond radius, the length that is
half the distance between the nuclei of two
bonded atoms. The bond radius can change slightly
depending on what atoms are involved.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Atomic Radius Increases as You Move Down a Group
As you proceed from one element down to the next
in a group, another principal energy level is
filled. The addition of another level of
electrons increases the size, or atomic radius,
of an atom.
Atomic Radius Decreases as You Move Across a
Period
As you move from left to right across a period,
each atom has one more proton and one more
electron than the atom before it has. As a
result, electron shielding does not play a role
as you move across a period. Therefore, as the
nuclear charge increases across a period, the
effective nuclear charge acting on the outer
electrons also increases. This increasing
nuclear charge pulls the outermost electrons
closer and closer to the nucleus and thus
reduces the size of the atom.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Electronegativity
Atoms often bond to one another to form a
compound. These bonds can involve the sharing of
valence electrons. Not all atoms in a
compound share electrons equally. Knowing how
strongly each atom attracts bonding electrons can
help explain the physical and chemical properties
of the compound. Linus Pauling, one of
Americas most famous chemists, made a scale of
numerical values that reflect how much an atom in
a molecule attracts electrons, called
electronegativity values. Chemical bonding that
comes from a sharing of electrons can be thought
of as a tug of war. The atom with the higher
electronegativity will pull on the electrons more
strongly than the other atom will.
Fluorine is the element whose atoms most
strongly attract shared electrons in a compound.
Pauling arbitrarily gave fluorine an
electronegativity value of 4.0.Values for the
other elements were calculated in relation to
this value.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Electronegativity Decreases as You Move Down a
Group
Metals tend to have small electronegativity
because they would rather lose an electron than
gain electrons to fulfill the octet rule.
As you go down a group the distance between the
valence electrons and the nucleus increase,
weakening the grip the nucleus has on it
valence electrons due to electron shielding.
Electronegativity Increases as You Move Across a
Period
Nonmetals tend to have large electronegativity
because they would rather gain an electron than
lose electrons to fulfill the octet rule.
As you go across a period the distance between
the valence electrons and the nucleus decreases,
strengthening the grip the nucleus has on it
valence electrons.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Periodic Trends in Ionic Size
Recall that atoms form ions by either losing or
gaining electrons. Like atomic size, ionic size
has periodic trends. As you proceed down a
group, the outermost electrons in ions are in
higher energy levels. Therefore, just as atomic
radius increases as you move down a group,
usually the ionic radius increases as well. These
trends hold for both positive and negative ions.
Metals tend to lose one or more electrons and
form a positive ion. As you move across a period,
the ionic radii of metal cations tend to
decrease because of the increasing nuclear
charge. As you come to the nonmetal elements in a
period, their atoms tend to gain electrons and
form negative ions. The diagram on the next slide
shows that as you proceed through the anions on
the right of a period, ionic radii still tend to
decrease because of the anions increasing
nuclear charge.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Summary of Ionic Radii
Atoms that form anions will increase in ionic
radius, due to the nuclear charge becoming less
than the total electron cloud charge.
Atoms that form cations will decrease in ionic
radius due to the nuclear charge becoming greater
than the electron cloud charge.
All atoms forming anions or cations will increase
ionic radius when moving down a group because of
the addition of a principle energy level.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Other information about the periodic table.
All elements with an atomic number greater than
83 have isotopes that are not stable and will
undergo nuclear decay. (They are radioactive)
Chemistry Reference tables lists,
electronegativity and ionizations energies and
atomic radii.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Which element forms an ion that is larger than
its atom? 1. Cl 2. Ca 3. Li 4. Mg
Which atom has a radius larger than the radius of
its ion? 1. Cl 2. Ca 3. S 4. Se
Which of the following elements in Period 3 has
the greatest metallic character? 1. Ar 2. Si 3.
Mg 4. S
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Which period of the Periodic Table contains more
metallic elements than nonmetallic elements? 1.
Period 1 2. Period 2 3. Period 3 4. Period 4
As the elements in Period 3 are considered from
left to right, they tend to 1. lose electrons
more readily and increase in metallic character
2. lose electrons more readily and increase in
nonmetallic character 3. gain electrons more
readily and increase in metallic character 4.
gain electrons more readily and increase in
nonmetallic character
Compared to the nonmetals in Period 2, the metals
in Period 2 generally have larger 1. ionization
energies 2. electronegativities 3. atomic radii
4. atomic numbers
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Which of the following Group 2 elements has the
lowest first ionization energy? 1. Be 2. Mg 3.
Ca 4. Ba
Which of the following atoms has the greatest
tendency to attract electrons? 1. barium 2.
beryllium 3. boron 4. bromine
In which shell are the valence electrons of the
elements in Period 2 found?   1. 1   2. 2   3.
3   4. 4
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
Which of the following ions has the smallest
radius? 1. F-  2. Cl-  3. K 4. Ca2
The ability of carbon to attract electrons is 1.
greater than that of nitrogen, but less than that
of oxygen 2. less than that of nitrogen, but
greater than that of oxygen 3. greater than that
of nitrogen and oxygen 4. less than that of
nitrogen and oxygen
Which trends appear as the elements in Period 3
are considered from left to right? 1. Metallic
character decreases, and electronegativity
decreases. 2. Metallic character decreases, and
electronegativity increases. 3. Metallic
character increases, and electronegativity
decreases. 4. Metallic character increases, and
electronegativity increases.
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
As the atoms in Period 3 of the Periodic Table
are considered from left to right, the atoms
generally show 1. an increase in radius and an
increase in ionization energy 2. an increase in
radius and a decrease in ionization energy 3. a
decrease in radius and an increase in ionization
energy 4. a decrease in radius and a decrease in
ionization energy
Which Group 16 element has only unstable
isotopes? 1. Po 2. Te 3. Se 4. S
How much energy is required to remove the most
loosely bound electron from a neutral atom of
carbon in the gaseous phase?  1. 801 kJ/mol  2.
1086 kJ/mol  3. 1251 kJ/mol  4. 1000 kJ/mol
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Chapter 4 The Periodic Table
Section 3 - Trends in the Periodic Table
THE END
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Review
How is periodic table arranged
In order of atomic number not mass as the
original periodic tables were.
Define
Groups
Go up and down on periodic table
Periods
Go across the periodic table
Alkali Metals
Group 1
Alkaline Earth Metals
Group 2
Group 17
Halogens
Noble Gases
Group 18
Transition Elements
Groups 3-12
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Left
Right
Yes
No
Yes
No
Yes
No
Yes
No
Lose
Gain
Mercury
Bromine
Francium
Fluorine
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Ionization Energy
Energy required to remove one valence electron
Electronegativity
Willingness to gain an electron
Atomic Radii
Size of an atom
Ionic Radii
Size of an atom after it has gained/lost an
electron(s)
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Decrease due to increase in principle energy
levels
Increase due to increase in nuclear charge
Increase due to increase in nuclear charge
Decrease due to larger radii
Increase due to increase in principle energy
levels
Decrease due to increase in nuclear charge
Increase due to increase in principle energy
levels. Does not matter if gaining or losing
electrons
Varies- If gaining electrons it will increase if
losing electrons it will decrease
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