Periodic Properties of the Elements

Introduction

The periodic table is a foundational tool in chemistry, organizing elements according to recurring chemical and physical properties. Understanding periodic trends allows chemists to predict element behavior and reactivity.

Development of the Periodic Table

Historical Background

• Dmitri Mendeleev and Lothar Meyer independently proposed that elements should be grouped by their properties.

• Mendeleev used chemical properties to organize the table and predicted the existence and properties of undiscovered elements.

• Example: Germanium (Ge) was predicted as "eka-silicon" before its discovery.

Table: Comparison of Predicted and Observed Properties of Germanium

Atomic Number

Definition and Importance

• Mendeleev’s table was originally based on atomic mass, but the modern table is organized by atomic number.

• Atomic number (Z) is the number of protons in the nucleus of an atom.

• Henry Moseley established the concept of atomic number experimentally using X-rays.

Periodicity

Definition and Key Properties

• Periodicity refers to the repetitive pattern of properties as a function of atomic number.

• Main periodic properties discussed:

  • Sizes of atoms and ions

  • Ionization energy

  • Electron affinity

  • Select group chemical property trends

Effective Nuclear Charge (Zeff)

Concept and Calculation

• Effective nuclear charge (Zeff) is the net positive charge experienced by an electron in a multi-electron atom.

• Electrons are attracted to the nucleus but are also repelled by other electrons, which screen or shield the nuclear charge.

• The effective nuclear charge is less than the actual nuclear charge due to this shielding effect.

• General formula: where is the atomic number and is the screening constant.

Coulomb’s Law

Electrostatic Forces in Atoms

• Coulomb’s Law describes the force between two charged particles:

• Where and are the charges, is the distance between them, and is a proportionality constant.

• Increasing the magnitude of charges or decreasing the distance increases the force.

Hydrogen vs. Many-Electron Atoms

Screening and Shielding Effects

• In hydrogen, the electron feels the full nuclear charge.

• In multi-electron atoms, electrons are screened from the nucleus by other electrons, reducing the effective nuclear charge.

• This screening effect makes the attraction less than it would be in the absence of other electrons.

Trends in Effective Nuclear Charge

Across a Period and Down a Group

• Across a period: Zeff increases as the number of protons increases but shielding does not increase as much.

• Down a group: Zeff increases slightly, but the effect is less pronounced due to increased shielding from additional electron shells.

Atomic Size (Atomic Radius)

Factors Affecting Atomic Radius

• Principal quantum number (): Higher means larger atomic radius.

• Nuclear charge (): Higher pulls electrons closer, reducing atomic radius.

Trends

• Across a period: Atomic radius decreases due to increasing Zeff.

• Down a group: Atomic radius increases due to higher principal quantum number.

Sizes of Ions (Ionic Radius)

Factors Affecting Ionic Radius

• Nuclear charge

• Number of electrons

• Orbitals/energy levels occupied by valence electrons

Comparing Ionic Sizes

• Cations (positive ions) are smaller than their parent atoms due to loss of electrons and increased Zeff.

• Anions (negative ions) are larger than their parent atoms due to gain of electrons and decreased Zeff.

• Ions with the same charge can be compared directly; ions with the same number of electrons (isoelectronic series) can be compared by nuclear charge.

Ionization Energy

Definition and Trends

• Ionization energy (In) is the energy required to remove an electron from a gaseous atom or ion.

• First ionization energy ():

• Second ionization energy ():

• Successive ionization energies increase:

• Large jump in ionization energy after removal of all valence electrons (core electrons are much harder to remove).

Trends

• Across a period: Ionization energy increases due to higher Zeff.

• Down a group: Ionization energy decreases due to increased atomic size and lower Zeff felt by valence electrons.

Electron Configuration of Ions

Formation of Cations and Anions

• For cations: Electrons are removed from the highest principal quantum number () orbital first.

• For anions: Electrons are added to the lowest available energy orbital.

Electron Affinity

Definition and Trends

• Electron affinity (EA) is the energy change when an electron is added to a gaseous atom to form an anion.

• Example:

• Electron affinity is typically negative (energy released), indicating attraction between atom and electron.

• Noble gases have positive or near-zero electron affinity (energy must be supplied).

Trends

• Across a period: Electron affinity becomes more negative (stronger attraction for electrons).

• Down a group: Electron affinity becomes less negative (weaker attraction for electrons).

Table: Electron Affinity Values for Selected Groups

Additional info: Trends and values are generalized; actual values may vary slightly depending on measurement methods.