Atomic Structure
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Atoms are mostly empty space with a small, dense nucleus at the center.
The nucleus contains protons (positive charge) and neutrons (neutral charge).
Electrons (negative charge) orbit the nucleus in specific energy levels or shells.
Subatomic Particles
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Protons: Positive charge (+1), relative mass of 1.
Neutrons: Neutral charge (0), relative mass of 1.
Electrons: Negative charge (-1), relative mass of 1/1836 (very small).
Atomic Number and Mass Number
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Atomic Number: The number of protons in the nucleus of an atom. It also equals the
number of electrons in a neutral atom.
Mass Number: The total number of protons and neutrons in the nucleus of an atom.
Number of Neutrons: Mass Number - Atomic Number.
Mass and Charge Distribution
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Most of the atom's mass is concentrated in the nucleus due to the presence of protons
and neutrons.
The nucleus has a positive charge due to the protons.
Electrons orbit the nucleus, contributing very little to the overall mass. They create a
"cloud" of negative charge around the nucleus.
The electrostatic attraction between the positively charged nucleus and the negatively
charged electrons holds the atom together.
 Subatomic Particles and their Behavior in Electric Fields:
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Electrons: Negatively charged, deflected towards the positive plate in an electric field.
Have a small mass due to their significant deflection.
Protons: Positively charged, deflected towards the negative plate in an electric field.
Deflected less than electrons, indicating they have a greater mass.
Neutrons: Neutral, not deflected by electric fields.
 Determining the Subatomic Structure of Atoms and Ions:
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Atoms: Neutral; number of protons = number of electrons.
Ions: Formed when atoms gain or lose electrons.
o Positive ions (cations): Lose electrons, resulting in fewer electrons than protons.
o Negative ions (anions): Gain electrons, resulting in more electrons than protons.
Protons: Number of protons = atomic number.
Neutrons: Number of neutrons = mass number - number of protons.
 Atomic Radius:
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The measure of the size of an atom.
Defined as half the distance between the nuclei of two covalently bonded atoms of the
same type.
 Atomic Radius Trends:
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Across a period: Generally decreases due to increasing nuclear charge (more protons)
pulling the electrons closer to the nucleus.
Down a group: Generally increases due to the addition of electron shells.
Electron Shell Theory
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Atomic radii decrease across a period: As you move across a period in the periodic
table, the atomic number increases, meaning the number of protons and electrons
increases. These extra electrons are added to the same principal energy level or shell. The
increased nuclear charge (more protons) pulls the electrons closer to the nucleus,
resulting in a smaller atomic radius.
Atomic radii increase down a group: As you move down a group in the periodic table,
new energy levels (shells) are added. The outer electrons are further away from the
nucleus, leading to a larger atomic radius. Additionally, the inner electrons shield the
outer electrons from the full attractive force of the nucleus, further increasing the atomic
radius.
Ionic Radius
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Ionic Radius: The measure of the size of an ion.
Ions: Charged atoms (formed by gaining or losing electrons).
Anions (Negative Ions): Gain electrons.
o Increase in size due to electron repulsion and reduced nuclear attraction.
Cations (Positive Ions): Lose electrons.
o Decrease in size due to increased nuclear attraction per electron.
Electron Shell Theory
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Negative Ions: Gain electrons, increasing electron-electron repulsion. This pushes the
electrons further from the nucleus, making the ion larger.
Positive Ions: Lose electrons, reducing electron repulsion.
This allows the remaining electrons to be pulled closer to the nucleus, making the ion
smaller.
Isotopes
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Definition: Atoms of the same element with the same number of protons but different
numbers of neutrons.
Symbol: Chemical symbol followed by a dash and the mass number (e.g., ¹²C, ¹³C).
Chemical Properties:
o Isotopes of the same element have the same chemical properties.
o This is because they have the same number of electrons, which determine
chemical behavior.
Physical Properties:
o Isotopes have slightly different physical properties due to their mass difference
(e.g., small differences in mass and density).
Calculating Abundance of Isotopes:
o The average atomic mass of an element is calculated using the weighted average
of the masses of its isotopes and their relative abundances.
o Formula: (M1 * x) + (M2 * (100-x)) = M(E)
 M1 and M2 are the masses of the isotopes.
 x is the relative abundance of the first isotope.
 M(E) is the average atomic mass of the element.
1. Electronic Shells
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Electrons are arranged around the nucleus in principal energy levels (or shells).
Principal quantum number (n): Used to number the energy levels.
o Lower n means closer to the nucleus and lower energy.
o Higher n means farther from the nucleus and higher energy.
Electron capacity: Each shell can hold a specific number of electrons:
o n = 1: 2 electrons
o n = 2: 8 electrons
o n = 3: 18 electrons
o n = 4: 32 electrons
2. Subshells
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Principal quantum shells are further divided into subshells.
Subshell designations: Represented by letters s, p, d, and f.
3. Orbitals
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Subshells contain one or more atomic orbitals.
Each orbital can hold a maximum of two electrons.
Number of orbitals in each subshell:
o s: one orbital (2 electrons)
o p: three orbitals (6 electrons)
o d: five orbitals (10 electrons)
o f: seven orbitals (14 electrons)
4. Shape of Orbitals
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Orbitals represent the regions of space around the nucleus where there is the
highest probability of finding an electron.
s-orbitals: Spherical in shape.
p-orbitals: Dumbbell-shaped, oriented along the x, y, and z axes.
d-orbitals: More complex shapes, including cloverleaf and doughnut-like shapes.
f-orbitals: Even more complex shapes.
Ground State
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Most stable electronic configuration of an atom.
Lowest energy state for electrons in an atom.
Achieved by filling the lower energy subshells first.
Order of filling subshells: Generally follows the Aufbau principle (increasing energy
order), but there are exceptions at higher energy levels (n = 3 and above).
Ionization Energy
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Definition: The energy required to remove one mole of electrons from one mole of
gaseous atoms of an element to form one mole of gaseous 1+ ions.
Measured under standard conditions.
Symbol: I or IE
Units: kJ mol⁻¹
Example:
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First Ionization Energy of Ca: Ca(g) → Ca⁺(g) + e⁻
Second Ionization Energy of Ca: Ca⁺(g) → Ca²⁺(g) + e⁻
Successive Ionization Energy
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Successive Ionization Energy refers to the energy required to remove additional
electrons from an ion.
Increases with each subsequent electron removal.
Reason:
o As electrons are removed, the positive charge on the ion increases.
o This leads to a stronger electrostatic attraction between the nucleus and the
remaining electrons.
o More energy is required to overcome this increased attraction and remove each
subsequent electron.
Factors Affecting Ionization Energy
1. Nuclear Charge:
o As atomic number increases, the positive nuclear charge also increases.
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Result: Stronger attraction between the nucleus and electrons, leading to higher
ionization energy across a period.
2. Distance from the Nucleus (Atomic Radius):
o Across a period: Atomic radius decreases due to increased nuclear charge
pulling electrons closer. This results in higher ionization energy.
o Down a group: Atomic radius increases due to the addition of electron shells.
This results in lower ionization energy.
3. Electron Shielding:
o Inner electrons shield outer electrons from the full attractive force of the nucleus.
o More inner shells: Greater shielding effect, making it easier to remove outer
electrons.
o Result: Ionization energy decreases down a group.
4. Electron-Electron Repulsion:
o Electrons in the same orbital repel each other.
o This repulsion makes it slightly easier to remove an electron from a filled or
partially filled subshell.
o Result: Can slightly decrease the first ionization energy in some cases.
General Ionization Energy Trends:
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Across a period: Generally increases due to increasing nuclear charge and decreasing
atomic radius.
Down a group: Generally decreases due to increasing atomic radius and increased
shielding.
Additional Points:
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Exceptions to the trends: There are some exceptions to these general trends, especially
for elements with full or half-filled subshells.
Second and subsequent ionization energies: Always increase due to the increasing
positive charge on the ion.