Chapter 3
Periodic Table and Periodicity of Properties
Topic 1
Introduction to the PeriodicTable
The Periodic Table is a fundamental tool in chemistry,
organizing the known chemical elements based on their atomic number and
recurring properties. Its development was a significant milestone in the
history of science, stemming from early attempts to classify elements in a
systematic manner.
Early Attempts at
Classification
In the 19th century, chemists
sought to arrange elements in a meaningful way, leading to the discovery of
periodic relationships. Notable early attempts include:
·
Dobereiner's
Triads: Dobereiner observed that
certain groups of three elements, called triads, had atomic masses that
followed a pattern. The atomic mass of the middle element was approximately the
average of the other two.
·
Newlands'
Octaves: Newlands noticed a repeating
pattern in the properties of elements when arranged in order of increasing
atomic mass, similar to musical notes. He proposed the "law of
octaves."
Mendeleev's Periodic Table
The most influential early
periodic table was created by Dmitri Mendeleev in 1869. He arranged
known elements based on their atomic masses and chemical properties, leading to
the formulation of the Periodic Law:
·
Periodic
Law: The properties of elements are
a periodic function of their atomic masses.
Mendeleev's table was
groundbreaking, as it predicted the properties of undiscovered elements and
allowed for the correction of inaccurate atomic masses.
Modern Periodic Table
With the discovery of atomic
number by Henry Moseley in 1913, the Periodic Law was refined to state that
the properties of elements are a periodic function of their atomic numbers.
This led to the modern periodic table, which is organized based on increasing
atomic number.
Key features of the modern
periodic table:
·
Periods: Horizontal rows representing elements with
the same number of electron shells.
·
Groups: Vertical columns representing elements with
similar chemical properties due to having the same number of valence electrons.
·
Blocks: Sections of the table based on the type of
orbital that receives the last electron (s, p, d, or f).
The periodic table is a
valuable tool for understanding the relationships between elements, predicting
their properties, and understanding chemical reactions. It provides a visual
representation of the underlying order in the structure of matter.
Topic 2
Periods and Groups in the
Periodic Table
Periods
·
Short
Period (Period 1): Contains only two
elements: hydrogen and helium.
·
Normal
Periods (Periods 2 and 3):
Each contains eight elements. Examples:
o Period 2: Lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine,
neon.
o Period 3: Sodium, magnesium, aluminum, silicon, phosphorus, sulfur,
chlorine, argon.
·
Long
Periods (Periods 4 and 5):
Each contains eighteen elements.
·
Very
Long Periods (Periods 6 and 7): These periods are longer due to the presence of the Lanthanides (elements 57-71) and Actinides (elements 89-103), which are typically placed
separately at the bottom of the periodic table to maintain a manageable format.
Groups
·
Group
1 (Alkali Metals): Hydrogen, lithium,
sodium, potassium, rubidium, cesium, francium. These elements have one valence
electron.
·
Group
2 (Alkaline Earth Metals):
Beryllium, magnesium, calcium, strontium, barium, radium. These elements have
two valence electrons.
·
Groups
3-12 (Transition Metals):
These elements are characterized by partially filled d orbitals in their outer
electron shells. Examples include iron, copper, zinc, and gold.
·
Groups
13-17 (Representative Elements): These elements have varying numbers of valence electrons and
exhibit a wide range of properties. Examples include boron, carbon, nitrogen,
oxygen, fluorine, aluminum, silicon, phosphorus, sulfur, chlorine.
·
Group
18 (Noble Gases): Helium, neon, argon,
krypton, xenon, radon. These elements have completely filled outer electron
shells and are generally unreactive.
Key Points:
·
The number of elements
in a period is determined by the maximum number of electrons that can occupy a
particular energy level (valence shell).
·
Elements within a
group share similar chemical properties due to having the same number of valence
electrons.
·
The placement of
Lanthanides and Actinides at the bottom of the periodic table is a convention
to maintain a compact format.
Example:
·
Period
3: The elements in this
period have 3 electron shells. The number of valence electrons increases from
left to right, from sodium (1 valence electron) to argon (8 valence electrons).
·
Group
1: All alkali metals
have 1 valence electron, giving them similar properties such as reactivity with
water and formation of ionic compounds.
Topic 3
Periodic Trends: Atomic Size
and Shielding Effect
Atomic Size
Atomic size is a measure of the
average distance between the nucleus and the outermost electrons of an atom. It's generally measured in picometers (pm).
Trends in Atomic Size:
·
Across
a Period:
o Decreases: As you move from left to right across a period, the
number of protons in the nucleus increases, which increases the effective
nuclear charge. This stronger attraction pulls the outer electrons
closer to the nucleus, reducing atomic size.
·
Down
a Group:
o Increases: As you move down a group, a new energy level (shell) is added
to the atom. This increased distance between the nucleus and the outermost
electrons leads to a larger atomic size.
Shielding Effect
The shielding effect is the
reduction in the attractive force between the nucleus and the outermost
electrons due to the presence of inner electrons.
Factors Affecting Shielding
Effect:
·
Number
of Inner Electrons: Atoms with more inner
electrons have a greater shielding effect.
·
Orbital
Penetration: Electrons in s
orbitals penetrate closer to the nucleus than those in p, d, or f orbitals,
leading to a stronger shielding effect.
Trends in Shielding Effect:
·
Across
a Period:
o Shielding effect remains relatively constant: Within a period, the number of inner
electrons remains the same, so the shielding effect doesn't change
significantly.
·
Down
a Group:
o Increases: As you move down a group, the number of inner electrons
increases, leading to a stronger shielding effect.
Relationship Between Atomic
Size and Shielding Effect:
·
A stronger shielding effect
reduces the effective nuclear charge felt by the outermost electrons, making
them less tightly bound to the nucleus. This results in a larger atomic size.
Example:
·
Sodium
(Na) vs. Potassium (K):
Both sodium and potassium are in Group 1. Potassium has a larger atomic size than sodium because it has a
greater number of inner electrons, which shield the outer electron more
effectively, reducing the effective nuclear charge.
In summary: Atomic size decreases across a
period due to increased effective nuclear charge, while it increases down a
group due to the addition of new energy levels. The
shielding effect, influenced by the number of inner electrons and orbital
penetration, plays a significant role in determining atomic size.
Topic 4
Ionization energy is the energy required to remove an electron
from a neutral gaseous atom. It is typically measured in kilojoules per mole
(kJ/mol).
In simpler terms, it's the
energy needed to break the bond between an atom and its outermost electron.
This process creates a positively charged ion (cation).
First Ionization Energy:
·
The energy needed to
remove the first electron from a neutral atom.
·
For example, the first
ionization energy of sodium (Na) is +496 kJ/mol.
Successive Ionization Energies:
·
Atoms with multiple
valence electrons can have multiple ionization energies.
·
Removing a second
electron requires more energy than removing the first, and so on.
·
Across
a Period:
o Increases: As you move from left to right across a period, the atomic size
decreases, and the effective nuclear charge increases. This stronger
attraction between the nucleus and valence electrons makes it more difficult to
remove an electron, resulting in higher ionization energies.
·
Down
a Group:
o Decreases: As you move down a group, the atomic size increases, and the
shielding effect of inner electrons becomes stronger. This reduces the
effective nuclear charge felt by the valence electrons, making it easier to
remove them, resulting in lower ionization energies.
Example:
·
Group
1 elements (alkali metals):
Have low first ionization energies because they have only one valence electron,
which is relatively easy to remove.
·
Group
17 elements (halogens):
Have high first ionization energies because they have seven valence electrons,
making it difficult to remove an electron without disrupting the stable
configuration.
In summary: Ionization energy increases across a period
due to increased effective nuclear charge and decreases down a group due to
increased atomic size and shielding effect. These trends are important for
understanding the reactivity and chemical properties of elements.
 |
Topic 5
Electron affinity is a measure of an atom's ability to attract
an additional electron. It's defined as the energy change (either released or
absorbed) when an electron is added to a neutral gaseous atom to form a
negative ion (anion).
Key points:
·
Energy
Release or Absorption:
Electron affinity can be either positive or negative:
o Positive: If energy is absorbed when an electron is added, the electronaffinity is positive, indicating that the atom is less likely to accept an
electron.
o Negative: If energy is released when an electron is added, the electron
affinity is negative, indicating that the atom has a strong affinity for
electrons.
·
Trend
Across a Period:
o Increases: Generally, electron affinity increases from left to right across
a period. As atomic size decreases, the nucleus has a stronger attraction for
electrons, making it more likely to accept one.
·
Trend
Down a Group:
o Decreases: Electron affinity generally decreases from top to bottom within a
group. As atomic size increases, the shielding effect of inner electrons becomes
stronger, reducing the attraction between the nucleus and the incoming
electron.
Examples:
·
Halogens
(Group 17): Halogens typically have high
negative electron affinities, indicating a strong tendency to gain electrons
and form anions.
·
Noble
Gases (Group 18): Noble gases have relatively
low electron affinities, as their filled outer electron shells make them
relatively stable and less likely to accept additional electrons.
Factors affecting electronaffinity:
·
Atomic
size: Smaller atoms generally have
higher electron affinities due to stronger nuclear attraction.
·
Shielding
effect: The presence of inner
electrons can shield the nucleus from the incoming electron, reducing the
attraction.
·
Electron
configuration: Atoms with half-filled or
completely filled subshells may have slightly higher electron affinities due to
the stability associated with these configurations.
In summary: Electron affinity is a measure of an atom's
ability to attract an electron. It generally increases across a period and
decreases down a group, reflecting trends in atomic size and shielding effects.
Understanding electron affinity is important for understanding the formation of
ions and chemical bonding.
Topic no 6
Electronegativity is a measure of an atom's ability to attract
shared electrons towards itself in a covalent bond. It's a crucial property in
understanding the nature of chemical bonds, particularly covalent bonds.
·
Across
a Period:
o Increases: Electronegativity generally increases from left to right across a
period. As atomic size decreases and effective nuclear charge increases, the
nucleus has a stronger attraction for electrons, including shared electrons in
covalent bonds.
·
Down
a Group:
o Decreases: Electronegativity generally decreases from top to bottom within a
group. As atomic size increases, the shielding effect of inner electrons
weakens the attraction between the nucleus and shared electrons.
Examples:
·
Group
17 elements (halogens):
Halogens have high electronegativities, making them highly electronegative
elements. This means they strongly attract shared electrons in covalent bonds,
often forming polar covalent bonds or ionic bonds.
·
Group
1 elements (alkali metals):
Alkali metals have low electronegativities, making them electropositive
elements. They are less likely to attract shared electrons and often form ionic
bonds by losing electrons.
Factors affecting
electronegativity:
·
Atomic
size: Smaller atoms have higher
electronegativities due to stronger nuclear attraction.
·
Shielding
effect: The presence of inner
electrons can shield the nucleus from shared electrons, reducing the
attraction.
·
Electron
configuration: Atoms with half-filled or
completely filled subshells may have slightly higher electronegativities due to
the stability associated with these configurations.
In summary: Electronegativity is a measure of an atom's
ability to attract shared electrons in a covalent bond. It increases across a
period and decreases down a group, reflecting trends in atomic size and
shielding effects. Electronegativity is crucial for understanding the polarity
of covalent bonds and the nature of chemical interactions between atoms.
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