- Basic
Building Blocks: Elements serve as the fundamental building blocks of
matter. Understanding their properties and behavior is essential for
comprehending the properties and reactions of all substances in the
universe.
- Increasing
Number of Elements: Over time, the number of known elements has grown
significantly. In 1800, only 31 elements were known, but by 1865, this
number had more than doubled to 63. Today, we have identified 114
elements, with ongoing efforts to create new ones.
- Complexity
of Individual Study: Investigating each element and its countless
compounds individually is a daunting task due to the sheer number of
elements and the numerous possible combinations. This complexity makes it
impractical to study each element in isolation.
- Systematic
Organization: Scientists recognized the need for a systematic way to
organize their knowledge about elements. This led to the development of
the periodic table, which arranges elements in a structured manner,
simplifying the study of their properties and relationships.
- Grouping
by Similarities: The periodic table groups elements with similar
properties together. Elements within the same group often share common
characteristics, which makes it easier to understand their behavior and
predict their reactions.
- Predictive
Power: The periodic table's arrangement enables scientists to predict
the properties of unknown elements based on the trends and patterns
observed among neighboring elements. This predictive capability aids in
the discovery and study of new elements.
- Rationalizing
Chemical Facts: The periodic table not only organizes existing
knowledge about elements but also rationalizes known chemical facts. It
provides a framework for explaining why certain elements exhibit specific
behaviors and reactions.
- Encouraging
Further Study: By highlighting gaps and patterns in the periodic
table, it encourages scientists to explore and study elements that may be
missing or poorly understood. This drives further research and exploration
in the field of chemistry.
- Education
and Communication: The periodic table is an invaluable tool for
teaching and learning chemistry. It simplifies the subject by offering a
visual representation of the relationships between elements, making it
accessible to students and educators.
- Scientific
Advancements: The classification of elements through the periodic
table has played a crucial role in advancing the field of chemistry. It
has guided research, inspired new theories, and led to the development of
innovative technologies and materials.
DEVELOPMENT
OF THE PERIODIC
TABLE
Development of the Periodic Table and the contributions of
various scientists, including Johann Dobereiner, A.E.B. de Chancourtois, John
Alexander Newlands, Dmitri Mendeleev, and Lothar Meyer:
- Johann
Dobereiner (Early 1800s):
- Dobereiner
was the first to suggest the idea of trends among the properties of
elements.
- He
noticed similarities among the physical and chemical properties of groups
of three elements called "Triads."
- In
each Triad, the middle element had an atomic weight approximately halfway
between the other two.
- His
idea, known as the "Law of Triads," was dismissed as a
coincidence since it only applied to a few elements.
- DOBEREIGNER’S
TRIAD
|
ELEMENTS
|
ATOMIC
WEIGHT
|
ELEMENTS
|
ATOMIC
WEIGHT
|
ELEMENTS
|
ATOMIC
WEIGHTS
|
|
Li
|
7
|
Ca
|
40
|
Cl
|
35.5
|
|
Na
|
23
|
Sr
|
88
|
Br
|
80
|
|
K
|
39
|
Ba
|
137
|
I
|
127
|
- A.E.B.
de Chancourtois (1862):
- A
French geologist, he arranged known elements in order of increasing
atomic weights and created a cylindrical table to show the periodic
recurrence of properties.
- His
work did not gain much attention at the time.
- John
Alexander Newlands (1865):
- Newlands
proposed the "Law of Octaves," arranging elements in order of
increasing atomic weights.
- He
noticed that every eighth element had properties similar to the first
element, akin to musical octaves.
- Newlands'
law worked for elements only up to calcium and was not widely accepted
initially.
- However,
he was later awarded the Davy Medal in 1887 by the Royal Society, London,
for his work.
- NEWLAND’S OCTAVES
|
ELEMENT
|
Li
|
Be
|
B
|
C
|
N
|
O
|
F
|
|
At.
Wt.
|
7
|
9
|
11
|
12
|
14
|
16
|
19
|
|
ELEMENT
|
Na
|
Mg
|
Al
|
Si
|
P
|
S
|
Cl
|
|
At.
Wt.
|
23
|
24
|
27
|
29
|
31
|
32
|
35.5
|
|
ELEMENT
|
K
|
Ca
|
|
|
|
|
|
|
At.
Wt.
|
39
|
40
|
|
|
|
|
|
- Dmitri
Mendeleev and Lothar Meyer (1869):
- Working
independently, both Mendeleev and Meyer proposed that when elements are
arranged by increasing atomic weights, similarities in properties occur
at regular intervals.
- Meyer
plotted physical properties like atomic volume, melting point, and
boiling point against atomic weight and observed a periodically repeating
pattern.
- Unlike
Newlands, Meyer noted changes in the length of this repeating pattern.
- Mendeleev,
a Russian chemist, is generally credited with developing the Modern
Periodic Table.
- Mendeleev
arranged elements in horizontal rows (periods) and vertical columns
(groups) in order of increasing atomic weights.
- Elements
with similar properties occupied the same vertical column.
- Mendeleev's
system was more elaborate and relied on a broader range of physical and
chemical properties.
- He
occasionally ignored strict atomic weight order to maintain similarities
in properties within groups.
- Mendeleev
left gaps in the table for undiscovered elements and successfully
predicted the existence and properties of elements like gallium and
germanium.
- His
bold quantitative predictions and their eventual success made him and his
Periodic Table famous.
- Mendeleev's
Periodic Table (1905):
- Mendeleev's
Periodic Table, published in 1905, is considered a milestone in the
development of the Periodic Table.
Development of the Modern Periodic Law and the present form of the
Periodic Table:
- Background
of Early Periodic Table Development:
- Mendeleev
developed his Periodic Table without knowledge of the internal structure
of the atom.
- At
the start of the 20th century, significant advancements in subatomic
particle theories began to emerge.
- Henry
Moseley's Contribution (1913):
- English
physicist Henry Moseley observed regularities in the X-ray spectra of
elements.
- Plotting
the square root of the frequency of X-rays (ν) against the atomic number
(Z) resulted in a straight line, rather than a plot against atomic mass.
- Moseley's
findings established that the atomic number (number of protons or
electrons in a neutral atom) is a more fundamental property than atomic
mass.
- This
discovery led to the modification of Mendeleev's Periodic Law, giving
rise to the Modern Periodic Law:
- "The
physical and chemical properties of the elements are periodic functions
of their atomic numbers."
- Significance
of Atomic Number and Quantum Numbers:
- Atomic
number, being equivalent to the nuclear charge or the number of electrons
in a neutral atom, became crucial in understanding the periodicity of
elements.
- Quantum
numbers and electronic configurations play a pivotal role in explaining
the periodic variation in the properties of elements.
- Development
of Various Forms of Periodic Tables:
- Over
time, numerous forms of the Periodic Table have been devised.
- Some
emphasize chemical reactions and valence, while others stress the
electronic configuration of elements.
- The
"long form" of the Periodic Table is the most widely used
today.
- Elements
in the Modern Periodic Table:
- In
the Modern Periodic Table, horizontal rows are referred to as
"periods," and vertical columns as "groups" or
"families."
- Elements
with similar outer electronic configurations are grouped in vertical
columns.
- The
International Union of Pure and Applied Chemistry (IUPAC) recommends
numbering the groups from 1 to 18, replacing the older notations.
- Number
of Periods and Elements:
- There
are a total of seven periods in the Periodic Table.
- The
period number corresponds to the highest principal quantum number (n) of
the elements within that period.
- The
first period contains 2 elements, while the subsequent periods contain 8,
8, 18, 18, and 32 elements, respectively.
- The
seventh period is incomplete and, like the sixth period, would
theoretically contain a maximum of 32 elements based on quantum numbers.
- Lanthanoids
and Actinoids:
- In
the Modern Periodic Table, elements from both the sixth and seventh
periods, known as lanthanoids and actinoids, respectively, are placed in
separate panels at the bottom of the table.
Nomenclature of elements with atomic numbers
greater than 100:
- Traditional
Naming Privileges:
- Historically,
the privilege of naming newly discovered elements rested with the
discoverer or discoverers.
- The
suggested names for new elements were ratified by the International Union
of Pure and Applied Chemistry (IUPAC).
- Controversy
Surrounding Naming:
- New
elements with very high atomic numbers are highly unstable and often
exist in extremely small quantities, sometimes just a few atoms.
- The
synthesis and characterization of these elements require sophisticated
and costly equipment and laboratories.
- Due
to the competitive nature of scientific research, disputes can arise,
with different scientists and laboratories claiming credit for the
discovery of the same element.
- For
instance, both American and Soviet scientists claimed to have discovered
element 104, leading to naming conflicts (Rutherfordium vs.
Kurchatovium).
- IUPAC's
Recommendation for Systematic Nomenclature:
- To
address naming disputes and controversies, IUPAC made a recommendation
that, until a new element's discovery is fully proven and its name
officially recognized, a systematic nomenclature should be used.
- This
systematic nomenclature is derived directly from the element's atomic
number, utilizing numerical roots for zero and numbers one to nine.
- Numerical
Roots and Nomenclature for Elements with Z > 100:
- The
numerical roots for elements with atomic numbers above 100 are shown in
Table
Notation for IUPAC Nomenclature
of Elements
|
DIGIT
|
NAME
|
Abbreviation
|
DIGIT
|
NAME
|
Abbreviation
|
|
0
|
ni
|
n
|
5
|
pent
|
p
|
|
1
|
un
|
u
|
6
|
hex
|
h
|
|
2
|
bi
|
b
|
7
|
sept
|
s
|
|
3
|
tri
|
t
|
8
|
oct
|
o
|
|
4
|
quad
|
q
|
9
|
enn
|
e
|
- These
roots are combined in the order of the digits that make up the atomic
number.
- The
suffix "ium" is added to these roots to form a temporary name
for the new element.
- Temporary
Naming and Symbol:
- Initially,
the new element is given a temporary name, often represented by a symbol
consisting of three letters.
- Permanent
Naming Process:
- Later,
a permanent name and symbol are assigned through a formal vote involving
IUPAC representatives from different countries.
- The
permanent name might reflect the country or state where the element was
discovered or pay tribute to a notable scientist.
- Current
Status:
- As
of now, elements with atomic numbers up to 118 have been discovered.
- IUPAC
has officially announced the names of all these elements, following the
systematic nomenclature and permanent naming process.
Electronic configurations of elements and the
long form of the Periodic Table:
- Electronic
Configurations Defined by Quantum Numbers:
- An
electron in an atom is defined by a set of four quantum numbers.
- The
principal quantum number (n) determines the main energy level or shell in
which the electron is located.
- Electrons
are distributed into different subshells, often referred to as orbitals
(s, p, d, f) within an atom.
- The
arrangement of electrons into these orbitals is known as the electronic
configuration of the element.
- Periods
Reflect Valence Shell Energy Levels:
- In
the Periodic Table, each period corresponds to the value of n for the
outermost or valence shell.
- Successive
periods are associated with the filling of the next higher principal
energy level (n = 1, n = 2, and so on).
- Number
of Elements in Each Period:
- The
number of elements in each period is twice the number of atomic orbitals
available in the energy level being filled.
- The
first period (n = 1) begins with the filling of the lowest level (1s) and
contains two elements: hydrogen (1s1) and helium (1s2).
- The
second period (n = 2) starts with lithium, and the third electron enters
the 2s orbital, followed by the filling of 2p orbitals, resulting in 8
elements.
- The
third period (n = 3) starts at sodium, and the added electron enters a 3s
orbital. The filling of 3s and 3p orbitals gives 8 elements.
- The
fourth period (n = 4) starts at potassium, and before the 4p orbital is filled,
the filling of 3d orbitals becomes energetically favorable. This leads to
the 3d transition series starting at scandium (Z = 21) and ending at zinc
(Z = 30). The fourth period contains 18 elements.
- The
fifth period (n = 5) is similar to the fourth, with the addition of the
4d transition series, beginning at yttrium (Z = 39). The period ends at
xenon, with 18 elements.
- The
sixth period (n = 6) consists of 32 elements and includes the filling of
6s, 4f, 5d, and 6p orbitals. The 4f-inner transition series starts at
cerium (Z = 58) and ends at lutetium (Z = 71), known as the lanthanoid
series.
- The
seventh period (n = 7) resembles the sixth, with the filling of 7s, 5f,
6d, and 7p orbitals. This period includes most of the man-made
radioactive elements and will end with the element having atomic number
118, belonging to the noble gas family.
- The
filling of 5f orbitals starts after actinium (Z = 89) and gives rise to
the 5f-inner transition series known as the actinoid series.
- Separate
Placement of Inner Transition Series:
- The
4f and 5f-inner transition series of elements (lanthanoids and actinoids)
are placed separately in the Periodic Table to maintain its structure and
classification principles, preserving elements with similar properties
within a single column.
ELECTRONIC
CONFIGURATIONS OF ELEMENTS
- Groupwise
Electronic Configurations:
- Elements
within the same vertical column or group in the Periodic Table share
similar valence shell electronic configurations.
- A
group is characterized by having the same number of electrons in the
outermost orbitals, leading to similar chemical properties among group
members.
- Example:
Group 1 Elements (Alkali Metals):
- Group
1 elements, known as alkali metals, all possess ns^1 valence shell
electronic configurations.
- The
ns1 configuration signifies that these elements have one
electron in their outermost s orbital.
- The
electronic configuration of alkali metals can be expressed as follows:
- Li
(Lithium): 1s22s1
- Na
(Sodium): 1s22s22p63s1
- K
(Potassium): 1s22s22p63s23p64s1
- Cs
(Cesium): 1s22s22p63s23p64s23d104p65s1
- These
elements exhibit similar chemical behavior due to their common valence
shell configuration, such as readily donating their outermost electron to
form cations with a +1 charge.
- Periodic
Dependence on Atomic Number:
- The
properties of elements show a periodic dependence on their atomic number,
which reflects the number of protons (and electrons) in the nucleus.
- This
periodicity in properties is a fundamental principle of the Periodic
Table and is not based on the relative atomic mass of elements.
- Theoretical
Foundation - Aufbau Principle:
- The
aufbau principle, based on the electronic configuration of atoms,
provides the theoretical foundation for the periodic classification of
elements.
- Chemical
Behavior and Groups/Families:
- Elements
within the same vertical column or group in the Periodic Table exhibit
similar chemical behavior.
- This
similarity arises because elements in a group have the same number and
distribution of electrons in their outermost orbitals.
- Classification
into Four Blocks:
- Elements
are classified into four blocks based on the types of atomic orbitals
that are being filled with electrons.
- The
four blocks are the s-block, p-block, d-block, and f-block.
- Exceptions
to Categorization:
- There
are two exceptions to this categorization:
- Helium,
strictly belonging to the s-block, is positioned in the p-block along
with other group 18 elements because it has a completely filled valence
shell (1s^2) and exhibits properties characteristic of noble gases.
- Hydrogen,
with only one s-electron, can be placed in group 1 (alkali metals). It
can also gain an electron to achieve a noble gas arrangement and behave
similarly to group 17 (halogen family) elements.
- Positioning
of Hydrogen and Helium:
- Hydrogen
is placed separately at the top of the Periodic Table due to its unique
properties and behavior.
- Helium
is also included in the p-block with other noble gases due to its filled
valence shell.
- Brief
Overview of Element Types:
- s-block
Elements: These elements are found in the s-block of the Periodic
Table and typically have their outermost electrons filling the
s-orbitals. They include alkali metals and alkaline earth metals.
- p-block
Elements: Elements in this block have their outermost electrons
filling p-orbitals. They encompass a wide range of elements, including nonmetals,
metalloids, and some metals.
- d-block
Elements: This block comprises transition metals, which have their
outermost electrons filling d-orbitals. Transition metals are known for
their characteristic properties, including variable oxidation states and
the ability to form colored compounds.
- f-block
Elements: Elements in this block include the lanthanides and
actinides, which have their outermost electrons filling f-orbitals. They
are often referred to as inner transition metals.
Characteristics of s-block elements:
- Classification
of s-Block Elements:
- The
s-block elements include Group 1 (alkali metals) and Group 2 (alkaline
earth metals) elements.
- They
share a common outermost electronic configuration, which is either ns1 for alkali metals
or ns2 for alkaline
earth metals, where "n" represents the principal quantum number
of the valence shell.
- Reactive
Metals:
- All
s-block elements are highly reactive metals with low ionization
enthalpies.
- They
readily lose their outermost electron(s) to form ions, resulting in a 1+
ion for alkali metals and a 2+ ion for alkaline earth metals.
- Their
reactivity and metallic character increase as you move down the group.
- Absence
in Pure Form in Nature:
- Due
to their high reactivity, s-block elements are rarely found in their
pure, uncombined form in nature.
- Instead,
they are typically found as compounds or minerals.
- Predominantly
Ionic Compounds:
- Compounds
formed by s-block elements, with the exception of those containing
lithium and beryllium, are predominantly ionic in nature.
- Ionic
compounds are characterized by the transfer of electrons from the metal
(s-block element) to the non-metal, resulting in the formation of ions
held together by electrostatic forces.
s-block elements, encompassing alkali metals and alkaline
earth metals, share common characteristics such as high reactivity, metallic
character, and the tendency to form ions. They readily lose their outermost
electrons to achieve stable electron configurations, resulting in the formation
of predominantly ionic compounds.
Characteristics of p-block elements:
- p-Block
Elements and Main Group Elements:
- The
p-Block Elements encompass elements belonging to Groups 13 to 18 in the
Periodic Table.
- Together
with the s-Block Elements, they are referred to as the Representative
Elements or Main Group Elements.
- Variation
in Outermost Electronic Configuration:
- Within
the p-block, the outermost electronic configuration varies from ns2 np1 to ns2 np6 in
each period.
- At
the end of each period in the p-block is a noble gas element with a
closed valence shell ns2
np6 configuration.
- Stability
of Noble Gases:
- Noble
gases, found at the end of each p-block period, have completely filled
valence shell orbitals with ns2
np6 electrons.
- These
noble gases exhibit very low chemical reactivity due to the stability of
their electron configuration.
- Halogens
and Chalcogens:
- Preceding
the noble gas family in the p-block are two chemically important groups
of nonmetals: the halogens (Group 17) and the chalcogens (Group 16).
- These
groups have highly negative electron gain enthalpies, making them readily
accept one or two electrons, respectively, to achieve a stable noble gas
configuration.
- Trends
in Chemical Behavior:
- As
you move from left to right across a period in the p-block, the
non-metallic character of elements increases.
- Conversely,
as you go down a group in the p-block, the metallic character of elements
increases.
The p-Block Elements, which include the Main Group Elements,
display a wide range of chemical behaviors. Elements in this block have varying
outermost electronic configurations, leading to different reactivity patterns.
Noble gases at the end of each period are extremely stable due to their fully
filled valence shell orbitals. Preceding them, the halogens and chalcogens are
notable for their electron acceptance tendencies and increasing non-metallic
character across a period.
Characteristics of d-block elements
(transition elements):
- Group
Range and Inner d Orbitals:
- The
d-Block Elements, also known as Transition Elements, include elements
from Group 3 to 12, located in the central part of the Periodic Table.
- These
elements are characterized by the filling of inner d orbitals by
electrons, hence the name d-Block Elements.
- General
Outer Electronic Configuration:
- Transition
elements share a common general outer electronic configuration of
(n-1)d^1-10ns^0-2, where "n" represents the principal quantum
number of the valence shell.
- Notably,
Pd has the electronic configuration 4d^10 5s^0.
- Metallic
Nature:
- All
transition elements are metals, characterized by properties such as
malleability, ductility, and electrical conductivity.
- Colorful
Ions and Variable Valence States:
- Transition
elements typically form colorful ions in their compounds due to the presence
of partially filled d orbitals.
- They
exhibit variable valence (oxidation) states, meaning they can exist in
different states of charge depending on the chemical reaction.
- Paramagnetism:
- Many
transition elements are paramagnetic, meaning they have unpaired
electrons in their electron configurations, making them attracted to a
magnetic field.
- Catalytic
Properties:
- Transition
elements are often used as catalysts in chemical reactions due to their
ability to provide an alternate reaction pathway with lower activation
energy.
- Exceptions:
Zn, Cd, and Hg:
- Zn
(zinc), Cd (cadmium), and Hg (mercury) have the electronic configuration
(n-1)d^10ns^2, which differs from other transition elements.
- These
three elements do not exhibit many of the typical properties of
transition elements, and they behave more like the s-block elements.
- Bridging
Role:
- Transition
metals serve as a bridge between the highly reactive metals of the
s-block elements and the less reactive elements of Groups 13 and 14.
- This
role earned them the name "Transition Elements."
The d-Block Elements, or Transition Elements, are
characterized by the filling of inner d orbitals, leading to their distinctive
properties. They are typically metals, known for forming colorful ions,
displaying variable valence states, and often acting as catalysts. Exceptions
include Zn, Cd, and Hg, which have a different electronic configuration and
behave more like s-block elements. Transition elements play a vital role in
chemistry due to their diverse properties and applications.
Characteristics of f-block elements (inner
transition elements):
- Lanthanoids
and Actinoids:
- The
f-Block Elements consist of two rows of elements located at the bottom of
the Periodic Table.
- The
first row is known as the Lanthanoids, ranging from Ce (Z = 58) to Lu (Z
= 71).
- The
second row is called the Actinoids, spanning from Th (Z = 90) to Lr (Z =
103).
- Outer
Electronic Configuration:
- These
elements are characterized by an outer electronic configuration of (n-2)f1-14
(n-1)d0–1 ns2.
- The
last electron added to each element is filled into the f-orbital, hence the
name f-Block Elements.
- Metals:
- All
f-Block Elements are metals, exhibiting typical metallic properties such
as electrical conductivity and malleability.
- Similar
Properties within Series:
- Within
each series (lanthanoids or actinoids), the properties of the elements
are quite similar due to the gradual filling of the f-orbitals.
- Complex
Chemistry of Actinoids:
- The
chemistry of actinoid elements, especially the early actinoids, is more
complex than that of the lanthanoids.
- Actinoids
can exhibit a wide range of oxidation states, making their chemistry
intricate and versatile.
- Radioactive
Nature:
- Many
of the actinoid elements are radioactive, posing challenges for handling
and study.
- These
elements often have short half-lives, and some can only be produced in extremely
small quantities through nuclear reactions.
- Transuranium
Elements:
- Elements
beyond uranium in the actinoid series are referred to as Transuranium
Elements.
- These
elements have atomic numbers greater than 92 (the atomic number of
uranium) and are generally synthetic, created through nuclear reactions.
The f-Block Elements, comprising the Lanthanoids and
Actinoids, are characterized by their unique electron configurations and
metallic nature. Elements within each series exhibit similar properties. The
chemistry of actinoid elements is notably complex due to their ability to adopt
various oxidation states. Many of these elements are radioactive, and those
beyond uranium are considered Transuranium Elements, created artificially in
small quantities.
METALS, NON-METALS AND
METALLOIDS
Metals:
- Abundance:
Metals make up more than 78% of all known elements and are primarily found
on the left side of the Periodic Table.
- State
at Room Temperature: Typically, metals are solids at room temperature,
with the exception of mercury (a liquid), and gallium and caesium (with
very low melting points).
- High
Melting and Boiling Points: Metals generally have high melting and
boiling points, contributing to their solid state at room temperature.
- Conductivity:
They are excellent conductors of heat and electricity, making them
essential materials in electrical and thermal applications.
- Malleability
and Ductility: Metals are malleable, meaning they can be flattened
into thin sheets by hammering, and ductile, allowing them to be drawn into
wires.
Non-Metals:
- Location:
Non-metals are predominantly found on the top right-hand side of the
Periodic Table.
- State
at Room Temperature: They are typically either solids or gases at room
temperature, with a few exceptions like boron and carbon.
- Low
Melting and Boiling Points: Non-metals generally have low melting and
boiling points, making them exist in various physical states.
- Poor
Conductors: They are poor conductors of both heat and electricity.
- Brittle
Nature: Most non-metallic solids are brittle and lack the malleability
and ductility observed in metals.
- Trend:
The metallic character of elements increases as one goes down a group,
while non-metallic character intensifies as one moves from left to right
across the Periodic Table.
Metalloids (Semi-metals):
- Location:
Metalloids are situated along the thick zig-zag line in the Periodic
Table, bordering the transition from metals to non-metals.
- Properties:
These elements, such as silicon, germanium, arsenic, antimony, and
tellurium, display properties that exhibit characteristics of both metals
and non-metals.
- Transitional
Nature: Metalloids have properties that make them intermediate between
metals and non-metals, contributing to their classification as
semi-metals.
Elements are broadly classified into Metals and Non-Metals based
on their properties. Metals are typically solid, have high melting points, and
are good conductors of heat and electricity. Non-Metals, on the other hand, are
often either solids or gases, possess lower melting points, and are poor
conductors. Metalloids, also known as Semi-metals, exhibit a mix of properties
from both categories and are situated along the boundary line between metals
and non-metals in the Periodic Table.
1. Patterns in Chemical Reactivity:
- Within
a Period (Horizontal Row):
- Reactivity
tends to be high in Group 1 metals (alkali metals) on the far left.
- It
decreases as you move towards the middle of the Periodic Table.
- It
then increases again to a maximum in Group 17 non-metals (halogens) on
the right.
- Within
a Group (Vertical Column):
- For
representative metals (e.g., alkali metals), reactivity increases as you
move down the group.
- For
non-metals (e.g., halogens), reactivity decreases as you move down the
group.
2. Explanation in Terms of Atomic Structure:
- Electron
Configuration and Energy Levels:
- Periodic
trends are explained by the distribution of electrons in an atom's energy
levels and orbitals.
- Elements
within a group have the same number of valence electrons (outermost
electrons), leading to similar chemical behavior.
- As
you move across a period, the number of valence electrons increases by
one from left to right.
- Effective
Nuclear Charge:
- The
effective nuclear charge, experienced by valence electrons, increases
across a period due to an increase in the number of protons in the nucleus.
- This
stronger attraction between the nucleus and valence electrons makes it
harder for elements to lose or gain electrons, affecting their
reactivity.
- Atomic
Size (Atomic Radius):
- Atomic
size generally decreases from left to right across a period.
- This
is because the increasing positive charge in the nucleus pulls electrons
closer to the nucleus, reducing the atomic radius.
- Ionization
Energy:
- Ionization
energy, the energy required to remove an electron from an atom, generally
increases across a period.
- As
atomic size decreases, electrons are held more tightly, making it harder
to remove them.
- Electronegativity:
- Electronegativity,
the ability of an atom to attract electrons in a chemical bond, increases
across a period.
- Smaller
atoms with higher effective nuclear charges have greater
electronegativity.
- Metallic
Character:
- Metallic
character decreases across a period.
- Metals
are on the left side of the Periodic Table, and as you move to the right,
elements become less metallic and more non-metallic in character.
- Group
Trends:
- In
a group, elements have the same valence electron configuration, resulting
in similar chemical properties.
- Reactivity
trends in a group are largely due to the increase in energy levels as you
move down the group, making valence electrons farther from the nucleus.
These periodic trends in properties can be explained by
considering the atomic structure, including electron configuration, effective
nuclear charge, and atomic size. These factors play a crucial role in
determining how elements react and behave within the Periodic Table.
Periodic trends in physical
properties of elements:
1. Atomic Radii:
- Across
a Period (Left to Right):
- Atomic
radii generally decrease as you move from left to right across a period.
- This
is due to the increase in effective nuclear charge, which pulls electrons
closer to the nucleus, resulting in smaller atomic sizes.
- Down
a Group (Top to Bottom):
- Atomic
radii generally increase as you move down a group.
- This
is because of the addition of new energy levels (shells) as you descend
the group, leading to larger atomic sizes.
2. Ionic Radii:
- Cations
(Positive Ions):
- Cations
are smaller than their parent atoms.
- When
an atom loses electrons to become a cation, the electron-electron
repulsion decreases, causing the remaining electrons to be pulled closer
to the nucleus.
- Anions
(Negative Ions):
- Anions
are larger than their parent atoms.
- When
an atom gains electrons to become an anion, the increased
electron-electron repulsion causes electrons to spread out, leading to larger
ionic sizes.
3. Ionization Enthalpy (Ionization Energy):
- Across
a Period (Left to Right):
- Ionization
enthalpy generally increases as you move from left to right across a
period.
- It
becomes harder to remove electrons because of the stronger effective nuclear
charge.
- Down
a Group (Top to Bottom):
- Ionization
enthalpy generally decreases as you move down a group.
- Electrons
are farther from the nucleus, experiencing less attraction, making them
easier to remove.
4. Electron Gain Enthalpy:
- Across
a Period (Left to Right):
- Electron
gain enthalpy tends to become more negative (more exothermic) as you move
from left to right across a period.
- Atoms
have a higher affinity for gaining electrons, except for noble gases.
- Down
a Group (Top to Bottom):
- Electron
gain enthalpy generally becomes less negative (less exothermic) as you
move down a group.
- Atoms
have a lower affinity for gaining electrons due to increased atomic size.
5. Electronegativity:
- Across
a Period (Left to Right):
- Electronegativity
generally increases as you move from left to right across a period.
- Atoms
become more electronegative due to the increasing effective nuclear
charge.
- Down
a Group (Top to Bottom):
- Electronegativity
generally decreases as you move down a group.
- Atoms
become less electronegative as the atomic size increases and valence
electrons are farther from the nucleus.
These trends in physical properties provide insights into
how the size of atoms, the ease of ionization, electron affinity, and
electronegativity change across the Periodic Table, influencing the chemical
behavior of elements.
PERIODIC TRENDS
IN ATOMIC RADIUS
1. Atomic Radius Estimation:
- Measuring
the size of an individual atom is challenging due to its tiny size (~1.2
Ã…).
- Atomic
size is estimated by measuring the distance between atoms in a covalent
molecule (covalent radius) or in a metallic crystal (metallic radius).
2. Covalent Radius for Non-Metals:
- Covalent
radius is half the distance between two atoms bound by a single covalent
bond in a molecule.
- Example:
Chlorine (Cl2) molecule has a bond distance of 198 pm, so the covalent
radius of chlorine is 99 pm.
3. Metallic Radius for Metals:
- Metallic
radius is half the internuclear distance between metal cores in a metallic
crystal.
- Example:
In solid copper, the distance between adjacent copper atoms is 256 pm, so
the metallic radius of copper is 128 pm.
4. Atomic Radii Measurement:
- Atomic
radii can be determined using various methods, including X-ray and
spectroscopic techniques.
5. Periodic Trends in Atomic Radius:
- Across
a Period (Left to Right):
- Atomic
radius generally decreases as you move from left to right across a
period.
- This
is because the effective nuclear charge increases due to a greater number
of protons, leading to stronger attraction of electrons to the nucleus.
- Down
a Group (Top to Bottom):
- Atomic
radius generally increases as you move down a group.
- This
is because electrons occupy higher energy levels (shells) as you descend
the group, causing them to be farther from the nucleus.
- Inner
energy levels shield outer electrons from the nucleus's pull,
contributing to larger atomic sizes.
6. Exception for Noble Gases:
- Noble
gases have extremely large atomic radii compared to other elements.
- Their
radii are not considered covalent radii but should be compared with van
der Waals radii of other elements.
These trends in atomic radius reflect how the size of atoms
changes across the Periodic Table due to variations in nuclear charge and
energy levels.
CONCEPTS OF IONIC RADIUS
AND ITS TRENDS
Concept of ionic radius and its
trends:
1. Ionic Radius Estimation:
- Ionic
radii are determined by measuring the distances between cations and anions
in ionic crystals.
- They
can be thought of as the sizes of ions formed when electrons are either
removed (cation) or added (anion) to an atom.
2. Trends in Ionic Radii:
- Cations
(Positive Ions):
- Cations
are smaller than their parent atoms.
- This
is because they have fewer electrons while maintaining the same nuclear
charge (protons).
- The
reduced electron-electron repulsion leads to a more compact ionic
structure.
- For
example, Na+ is smaller than neutral sodium (Na).
- Anions
(Negative Ions):
- Anions
are larger than their parent atoms.
- Adding
one or more electrons results in increased electron-electron repulsion
within the electron cloud.
- The
increased repulsion outweighs the effective nuclear charge, causing the
ion to expand.
- For
example, F- is larger than neutral fluorine (F).
3. Isoelectronic Species:
- Isoelectronic
species are atoms and ions that have the same number of electrons.
- Despite
having the same number of electrons, their ionic radii differ due to
varying nuclear charges.
- Greater
positive charge (cation) results in a smaller radius, while greater
negative charge (anion) leads to a larger radius.
- Example:
O2-, F-, Na+, and Mg2+ all
have 10 electrons, but their ionic radii vary due to differences in
nuclear charge.
These trends in ionic radii help explain the size variations
among cations and anions and are important in understanding the properties and
behavior of ionic compounds.
PERIODIC TRENDS
IN IONIZATION ENTHALPY
1. Ionization Enthalpy:
- Ionization
enthalpy is a quantitative measure of an element's tendency to lose
electrons.
- It
represents the energy required to remove an electron from an isolated
gaseous atom in its ground state.
- The
first ionization enthalpy refers to the energy needed to remove the first
electron from the atom.
- Ionization
enthalpies are always positive, indicating that energy is required to
remove electrons.
2. Trends in Ionization Enthalpy:
- Periodic
Trend Across Periods :
- Ionization
enthalpy generally increases as you move across a period (from left to
right).
- This
is due to the increasing effective nuclear charge as more protons are
added to the nucleus.
- Shielding
or screening by inner core electrons remains relatively constant across a
period, leading to stronger electron-nucleus attraction.
- As
a result, outermost electrons are held more tightly.
- Periodic
Trend Down Groups :
- Ionization
enthalpy generally decreases as you move down a group (from top to
bottom).
- This
is because the outermost electron is farther from the nucleus in lower
energy levels (higher principal quantum number, n).
- The
increased distance and increased shielding by inner electrons reduce the
effective nuclear charge on the outermost electron.
- Removal
of the outermost electron requires less energy.
3. Factors Influencing Ionization Enthalpy:
- Penetration
of Orbitals:
- When
comparing elements within the same principal quantum level, s-electrons
are more strongly attracted to the nucleus than p-electrons.
- For
example, boron (Z = 5) has a slightly lower first ionization enthalpy
than beryllium (Z = 4) because the electron removed from boron is a
p-electron, which is more shielded from the nucleus than the s-electron
in beryllium.
- Electron-Electron
Repulsion:
- Anomalies
in ionization enthalpy trends can occur due to electron-electron
repulsion within atomic orbitals.
- For
example, oxygen (Z = 8) has a lower first ionization enthalpy than
nitrogen (Z = 7) because oxygen's fourth 2p-electron experiences
increased repulsion since it must occupy the same 2p-orbital as the other
three 2p-electrons.
These trends in ionization enthalpy are influenced by both
the nuclear charge and the electron configuration of the elements, shedding
light on their reactivity and behavior in chemical reactions.
PERIODIC TRENDS
IN ELECTRON GAIN ENTHALPY
1. Electron Gain Enthalpy (ΔegH):
- Electron
gain enthalpy measures the energy change when an electron is added to a
neutral gaseous atom (X) to form a negative ion (X⁻).
- It
can be represented by the equation: X(g) + e⁻ → X⁻(g).
2. Nature of Electron Gain Enthalpy:
- Electron
gain enthalpy can be either endothermic or exothermic, depending on the
element.
- For
many elements, energy is released (exothermic) when an electron is added
to the atom, resulting in negative electron gain enthalpy.
- Group
17 elements (halogens) have highly negative electron gain enthalpies
because they can attain stable noble gas electronic configurations by
gaining an electron.
- Noble
gases have large positive electron gain enthalpies because adding an
electron leads to an unstable electronic configuration.
3. Variation Across Periods:
- Electron
gain enthalpies generally become more negative (exothermic) as you move
from left to right across a period.
- This
trend is due to the increasing effective nuclear charge (proton count)
across a period.
- Smaller
atoms have more negative electron gain enthalpies because the added
electron is closer to the positively charged nucleus.
4. Variation Down Groups:
- Electron
gain enthalpies tend to become less negative (endothermic) as you move
down a group.
- This
is because the atomic size increases down a group, and the added electron
is farther from the nucleus, requiring more energy.
- Electron
gain enthalpies of elements generally follow this trend.
5. Anomalies in Electron Gain Enthalpy:
- There
are some exceptions to the general trend, such as the electron gain
enthalpy of oxygen (O) or fluorine (F) being less negative than that of
the succeeding element in the same period.
- This
occurs because when an electron is added to O or F, it enters the smaller
n = 2 quantum level, causing significant repulsion from other electrons in
the same level.
- In
contrast, for the n = 3 quantum level (S or Cl), the added electron
occupies a larger region of space, resulting in less electron-electron
repulsion and more negative electron gain enthalpy.
These trends in electron gain enthalpy provide insights into
an element's ability to gain electrons and form anions, impacting its chemical
reactivity.
PERIODIC TRENDS
IN ELECTRONEGATIVITY
1. Electronegativity:
- Electronegativity
is a qualitative measure of an atom's ability in a chemical compound to
attract shared electrons towards itself.
- It
is not a directly measurable quantity but can be quantified using various
scales, with the Pauling scale being the most widely used.
2. Pauling Scale:
- Linus
Pauling assigned a value of 4.0 to fluorine, considering it to have the
greatest electron-attracting ability.
- Electronegativity
values for other elements are determined relative to this scale.
3. Variation Across Periods:
- Electronegativity
generally increases from left to right across a period (e.g., from lithium
to fluorine) in the periodic table.
- This
trend is explained by the increasing attraction between the outer
electrons and the nucleus as the atomic radius decreases across a period.
4. Variation Down Groups:
- Electronegativity
generally decreases as you move down a group (e.g., from fluorine to
astatine).
- This
trend is associated with the increase in atomic radii down a group,
resulting in weaker electron-nucleus attraction.
5. Relationship Between Electronegativity and Atomic
Radius:
- Electronegativity
and atomic radius exhibit an inverse relationship: as atomic radius
increases, electronegativity decreases, and vice versa.
6. Electronegativity and Non-Metallic Properties:
- Non-metallic
elements have a strong tendency to gain electrons, and this tendency is
directly related to electronegativity.
- The
increase in electronegativity across a period corresponds to an increase
in non-metallic properties (or a decrease in metallic properties) of
elements.
- Conversely,
the decrease in electronegativity down a group corresponds to a decrease
in non-metallic properties (or an increase in metallic properties) of
elements.
Understanding electronegativity helps predict chemical
behavior, especially in the formation of ionic and covalent compounds, and
provides insights into the relative tendencies of elements to attract
electrons.
PERIODIC TRENDS
IN CHEMICAL
PROPERTIES OF ELEMENTS
The periodic trends in chemical properties of elements,
focusing on valence states and anomalies in the second period elements:
1. Valence States:
- Elements
in the periodic table exhibit periodicity in the valence states they can
achieve.
- Valence
states are associated with the number of electrons an element can gain or
lose to attain a stable electron configuration.
- The
periodic table provides a framework for understanding the possible valence
states of elements within each group.
2. Diagonal Relationships:
- Diagonal
relationships refer to similarities in properties and behaviors between
elements in diagonally related positions in the periodic table.
- For
example, lithium (Li) and magnesium (Mg) show some similarities due to
their diagonal relationship.
3. Inert Pair Effect:
- The
inert pair effect is observed in heavier elements, particularly in the
p-block elements.
- It
refers to the tendency of some elements to preferentially lose or share
the s-electrons in their outermost electron shell, leaving the p-electrons
unaltered.
- This
results in the formation of ions with a lower oxidation state than
expected based on the group number.
4. Effects of Lanthanoid Contraction:
- Lanthanoid
contraction is a phenomenon observed in the transition metals,
specifically in the lanthanide series.
- It
involves a decrease in atomic and ionic radii across the lanthanide series
due to the poor shielding of inner electrons.
- This
contraction can affect the chemical properties and reactivity of elements
in this series.
5. Anomalies in Second Period Elements:
- Second
period elements from lithium (Li) to fluorine (F) exhibit several
anomalies in their properties compared to elements in other periods.
- These
anomalies are often attributed to the small size of second period atoms
and their high effective nuclear charge.
- Examples
include the relatively low ionization enthalpy of boron compared to
beryllium and the electron-electron repulsion in oxygen compared to
nitrogen.
These trends and phenomena in chemical properties provide
valuable insights into the behavior of elements and their compounds, helping
chemists understand and predict the outcomes of chemical reactions.
PERIODICITY OF
VALENCE OR OXIDATION STATES
1. Valence and Electronic Configurations:
- Valence
refers to the most characteristic property of elements, and it can be
understood in terms of their electronic configurations.
- For
representative elements, valence is often equal to the number of electrons
in the outermost orbitals (valence electrons).
- Alternatively,
valence can be calculated as eight minus the number of outermost
electrons.
2. Oxidation States:
- The
term "oxidation state" is frequently used interchangeably with
valence, especially in the context of chemical reactions and compounds.
- Oxidation
states are numerical values that represent the charge an atom would have
in a compound or ion.
- Oxidation
states can be positive, negative, or zero, depending on the electron
transfer in a chemical reaction.
3. Example: OF2 and Na2O:
- Consider
the compounds OF2 and Na2O.
- The
electronegativity order of the elements involved is F > O > Na.
- In
OF2, each fluorine atom tends to gain one electron to achieve a stable
noble gas configuration (F: 1s²2s²2p⁵ → F⁻: 1s²2s²2p⁶).
- Therefore,
each fluorine atom has an oxidation state of -1 in OF2.
- In
Na2O, sodium tends to lose one electron to achieve a noble gas
configuration (Na: 1s²2s²2p⁶3s¹ → Na⁺: 1s²2s²2p⁶).
- Oxygen,
on the other hand, tends to gain two electrons to attain a noble gas
configuration (O: 1s²2s²2p⁴ → O²⁻: 1s²2s²2p⁶).
- Therefore,
sodium has an oxidation state of +1, and oxygen has an oxidation state of
-2 in Na2O.
Understanding oxidation states and valence is crucial in
predicting and explaining chemical reactions and the formation of compounds. It
helps chemists balance chemical equations and determine the electron transfer
in reactions.
ANOMALOUS PROPERTIES
OF SECOND PERIOD
ELEMENTS
- Anomalous
Properties of First Group Elements:
- The
first element in each group, such as lithium in Group 1 and beryllium in
Group 2, and groups 13 to 17 (boron to fluorine), exhibits distinctive
differences compared to the other members within their respective groups.
- Covalent
Character in Lithium and Beryllium:
- Unlike
the other alkali metals, lithium and the other alkaline earth metal
beryllium tend to form compounds with significant covalent
characteristics. This is in contrast to the predominantly ionic compounds
formed by the other members of these groups.
- Similarity
with Second Elements in Adjacent Group:
- Interestingly,
the chemical behavior of lithium resembles that of the second element in
the following group, magnesium, and beryllium's behavior is akin to that
of aluminum, the second element in its group. This similarity is referred
to as a "diagonal relationship" in periodic properties.
- Reasons
for Anomalous Behavior:
- The
anomalous behavior of the first group element compared to its group
members is primarily attributed to several factors:
- Small
size: The first member has a smaller atomic size.
- Large
charge-to-radius ratio: It possesses a higher charge-to-radius ratio.
- High
electronegativity: These elements tend to be highly electronegative.
- Limited
Valence Orbitals:
- The
first member of each group has only four valence orbitals (2s and 2p)
available for bonding.
- Covalency
Limitation:
- Consequently,
the maximum covalency of the first group member is limited to 4. For
example, boron can form compounds like BF₄⁻, but it cannot expand its
valence shell beyond four pairs of electrons.
- Greater
Ability for π Bonds:
- Additionally,
the first member of the p-block elements shows a greater propensity to
form pÏ€ – pÏ€ multiple bonds with itself (e.g., C = C, C ≡ C, N = N, N ≡
N) and with other second-period elements (e.g., C = O, C = N, C ≡ N, N =
O). This ability diminishes in subsequent group members.
- Expanding
Valence Shell:
- In
contrast, the second member of each group, with nine valence orbitals
(3s, 3p, 3d), can expand its valence shell to accommodate more than four
pairs of electrons, resulting in a wider range of bonding possibilities
(e.g., aluminum can form AlF₆³⁻).
PERIODIC TRENDS
AND CHEMICAL REACTIVITY
- Fundamental
Properties and Periodicity:
- All
chemical and physical properties of elements are linked to their
electronic configurations. Understanding periodic trends in properties
such as atomic and ionic radii, ionization enthalpy, electron gain
enthalpy, and valence is essential.
- Periodic
Trends in Radii and Enthalpies:
- Across
a period, atomic and ionic radii generally decrease from left to right.
Consequently, ionization enthalpies tend to increase. Exceptions exist,
as mentioned earlier.
- Chemical
Reactivity and Enthalpies:
- Elements
on the far left and right ends of a period exhibit the highest chemical
reactivity. The leftmost element loses electrons to form cations, while
the rightmost element gains electrons to form anions. This behavior
relates to the metallic and non-metallic character of elements.
- Metallic
and Non-Metallic Character:
- Moving
from left to right across a period, metallic character decreases, and
non-metallic character increases. Alkali metals on the far left are
highly metallic, while halogens on the far right are non-metals.
- Reactivity
with Oxygen:
- Chemical
reactivity is evident in reactions with oxygen. Elements at the extremes
of a period readily form oxides. Leftmost elements produce basic oxides
(e.g., Na₂O), while rightmost elements produce acidic oxides (e.g.,
Cl₂O₇). Central elements produce amphoteric (e.g., Al₂O₃, As₂O₃) or
neutral (e.g., CO, NO, N₂O) oxides.
- Transition
Metals:
- Transition
metals (3d series) exhibit smaller changes in atomic radii across a
period compared to representative elements. Their ionization enthalpies
are intermediate, making them less electropositive than group 1 and 2
metals.
- Group
Trends:
- Within
a group, as atomic and ionic radii increase with atomic number,
ionization enthalpies generally decrease. This leads to an increase in
metallic character and a decrease in non-metallic character down the
group.
- Transition
Elements in Groups:
- In
contrast, transition elements show a reverse trend within groups,
primarily due to atomic size and ionization enthalpy considerations.
3.1 What is the basic theme of organisation in the periodic
table?
3.2 Which important property did Mendeleev use to classify
the elements in his periodic table and did he stick to that?
3.3 What is the basic difference in approach between the
Mendeleev’s Periodic Law and the Modern Periodic Law?
3.4 On the basis of quantum numbers, justify that the sixth
period of the periodic table should have 32 elements.
3.5 In terms of period and group where would you locate the
element with Z =114?
3.6 Write the atomic number of the element present in the
third period and seventeenth group of the periodic table.
3.7 Which element do you think would have been named by (i)
Lawrence Berkeley Laboratory (ii) Seaborg’s group?
3.8 Why do elements in the same group have similar physical
and chemical properties?
3.9 What does atomic radius and ionic radius really mean to
you?
3.10 How do atomic radius vary in a period and in a group?
How do you explain the variation?
3.1 Basic Theme of Organization in the Periodic Table:
- The
basic theme of organization in the periodic table is the arrangement of
chemical elements in a systematic and ordered manner based on their atomic
number, electron configuration, and recurring chemical properties.
Elements are grouped into periods (horizontal rows) and groups (vertical
columns) to emphasize similarities and trends in their properties.
3.2 Mendeleev's Use of Property in Classification:
- Dmitri
Mendeleev classified elements in his periodic table primarily based on
their atomic mass. He organized the elements in order of increasing atomic
mass and noticed that elements with similar chemical properties occurred
at regular intervals. However, Mendeleev did not strictly stick to atomic
mass; he adjusted the order of elements when it better fit their chemical
properties, sometimes even leaving gaps for undiscovered elements.
3.3 Difference between Mendeleev's and Modern Periodic
Law:
- Mendeleev's
Periodic Law was based on atomic mass and chemical properties, and he left
gaps for undiscovered elements. The key difference is that Mendeleev did
not have a theoretical explanation for his periodic table, while the
Modern Periodic Law is based on the atomic number (number of protons) and
is explained by the electronic configuration of elements. In the modern
table, elements are arranged in order of increasing atomic number, and
there are no gaps.
3.4 Quantum Numbers and the Sixth Period:
- According
to the quantum mechanical model, each period of the periodic table
corresponds to the filling of a new energy level or shell. The maximum
number of electrons that can occupy a given energy level is determined by
the formula 2n², where 'n' is the principal quantum number. For the sixth
period, 'n' is equal to 6, so the maximum number of elements that can be
accommodated is 2 × 6² = 72. Since there are already 18 elements in the
sixth period, 72 - 18 = 54 elements could potentially fill the sixth
period, resulting in a total of 18 + 54 = 72 elements in the sixth period.
3.5 Locating Element with Z = 114:
- An
element with Z = 114 would be located in the seventh period (row) of the
periodic table and in the fourteenth group (column), often referred to as
Group 14.
3.6 Atomic Number in the Third Period and Seventeenth
Group:
- The
element in the third period and seventeenth group of the periodic table
has an atomic number of 17. This element is chlorine (Cl).
3.7 Naming Elements by Research Institutions:
- (i)
The element with Z = 102, which is No. 102 on the periodic table, was
named "Nobelium" by the Lawrence Berkeley Laboratory.
- (ii)
Elements beyond No. 102 were discovered by Seaborg's group, and they are
known as transuranium elements. For example, element No. 106 was named
"Seaborgium" in honor of Glenn T. Seaborg, who played a
significant role in the discovery of these elements.
3.8 Similar Properties in the Same Group:
- Elements
in the same group have similar physical and chemical properties because
they have the same number of valence electrons. Valence electrons are
responsible for an element's chemical reactivity, and elements in the same
group have the same outer electron configuration, leading to similar
chemical behavior.
3.9 Atomic and Ionic Radius:
- Atomic
radius refers to the size of an atom, typically measured as the distance
from the nucleus to the outermost electron cloud. It reflects the size of
the electron cloud.
- Ionic
radius refers to the size of an ion, which can be larger or smaller than
the atomic radius depending on whether the ion has gained or lost
electrons.
3.10 Variation of Atomic Radius:
- In
a period (horizontal row), atomic radius generally decreases from left to
right. This is because, within a period, the number of protons in the
nucleus increases, leading to a greater nuclear charge pulling the
electrons closer to the nucleus.
- In
a group (vertical column), atomic radius generally increases as you move down
the group. This is because each new element in a group has an additional
energy level (shell) of electrons, making the electron cloud larger.
- The
variation in atomic radius can be explained by the balance between
increasing nuclear charge and increasing electron shells as you move
across periods and down groups in the periodic table.
