Periodic Properties and Trends

Periodic Properties of the Elements

  • Elements in the same group generally have similar chemical properties, including:

    • Reactivity

    • Flammability

    • Toxicity

  • Physical properties are not necessarily similar, including:

    • Color

    • Density

    • Appearance

    • Melting Point (MP) / Boiling Point (BP)

Development of the Periodic Table

  • Dmitri Mendeleev and Lothar Meyer independently concluded how elements should be grouped based on chemical properties.

  • Mendeleev predicted the discovery of germanium (known as eka-silicon) with an atomic weight between zinc and arsenic but with chemical properties similar to silicon.

Periodic Trends

  • Observed trends include:

    • Electron configuration

    • Effective Nuclear Charge

    • Sizes of atoms and ions

    • Ionization energy

    • Electron affinity

Electron Configurations

  • Definition: Electron configurations denote the specific orbitals that electrons occupy for a given atom or ion.

    • Ground state: Lowest energy state of an atom.

    • Excited state: A higher energy state where an electron is promoted to a higher energy level.

    • Aufbau Principle: Electrons fill orbitals starting from the lowest energy levels ascending to higher levels. The periodic table is used as a map to predict electron configurations.

Pauli Exclusion Principle

  • Definition: No two electrons in an atom can have the same set of four quantum numbers.

    • Consequently, each orbital can hold a maximum of two electrons, which must have opposite spins.

Applying the Pauli Exclusion Principle

  • Maximum number of electrons in sublevels based on the number of orbitals:

    • s sublevel: 1 orbital, holds 2 electrons.

    • p sublevel: 3 orbitals, holds 6 electrons.

    • d sublevel: 5 orbitals, holds 10 electrons.

    • f sublevel: 7 orbitals, holds 14 electrons.

Allowed Quantum Numbers

  • Quantum numbers for an electron in a 4s orbital could be:

    • n = 4

    • l = 0

    • m_l = 0

    • m_s = +1/2 (If one electron)

    • or m_s = -1/2 (If another electron)

  • For a 2p electron, possible quantum numbers include:

    • n = 2

    • l = 1

    • m_l = -1, 0, +1

    • m_s = +1/2 or -1/2

Determining Electron Configurations

  1. Determine the total number of electrons in the atom or ion.

  2. Start filling electrons from the ground state using the Aufbau principal.

    • Each orbital can hold two electrons, which must be spin paired (ms = +1/2 and -1/2).

  3. Stop filling when there are no electrons left.

Orbital Energies by Quantum Number

  • The energy order of orbitals with the principal quantum number (n):

    • The order of subshell energies is:

    • ns < np < nd < nf

    • Example for Ne configuration.

Periodic Table of Orbitals

  • The arrangement of orbitals is as follows:

    • 1s

    • 2s

    • 3s

    • 4s

    • 5s

    • 6s

    • 7s (continued with 3p, 4p, 4d, etc.)

Electron Configuration Examples

Example box diagrams for early elements

  • Lithium (Li):

    • Number of electrons: 3

    • Electron Configuration: 1s^2 2s^1

    • Orbital Diagram:

    • 1s (filled with two arrows),

    • 2s (half-filled with one arrow)

  • Beryllium (Be):

    • Number of electrons: 4

    • Electron Configuration: 1s^2 2s^2

    • Orbital Diagram:

    • 1s (filled),

    • 2s (filled)

  • Boron (B):

    • Number of electrons: 5

    • Electron Configuration: 1s^2 2s^2 2p^1

    • Orbital Diagram:

    • 1s (filled),

    • 2s (filled),

    • 2p (first two boxes half filled, one arrow each; third box empty)

Box diagrams vs. Written Electron Configurations

  • Condensed electron configurations for an element (e.g., Rubidium):

    • Closest noble gas before Rb:

    • [Kr] 4s^1

Valence Electrons

  • Defined as the electrons in the highest principal energy shell, while electrons in lower energy shells are core electrons.

  • The number of valence electrons influences an atom's chemical and physical behavior significantly.

Relationship of Electron Configuration to the Periodic Table

  • The main group number corresponds to the number of valence electrons.

  • The length of each block in the periodic table signifies the maximum number of electrons the sublevel can hold.

  • The period number indicates the principal energy level of the valence electrons.

Nuclear Charge

  • Definition: Nuclear charge (Z) is the net positive charge experienced by electrons in an atom due to the nucleus.

    • Each electron is attracted to the nucleus but also repelled by other electrons.

  • Effective nuclear charge (Z_eff) can be defined as:

    • Z_{eff} = Z - S

    • where:

    • Z = atomic number

    • S = screening (or shielding constant)

Shielding and Effective Nuclear Charge

Example with Lithium:

  • The effective nuclear charge felt by a valence (2s) electron:

  • Calculation: $Z_{eff} ≈ (3^+) - 2 = 1^+$

Atomic Radius

  • Definition: Bonding atomic radius is defined as half the distance between covalently bonded nuclei.

  • Trends in Size:

    • Atomic size decreases from left to right due to increased Z_eff.

    • Atomic size increases from top to bottom as n increases.

Sizes of Cations and Anions

  • Factors affecting ionic size include:

    • Nuclear charge

    • Number of electrons

    • Orbitals in which electrons reside

    • Cations: Smaller than their parent atoms, due to the removal of the outermost electron and reduction in repulsion.

    • Anions: Larger than their parent atoms, due to added electrons increasing repulsion.

  • Ionic size increases down a group due to increasing n.

Ionization Energy (IE)

  • Definition: Ionization energy is the energy required to remove an electron from a gaseous atom or ion's ground state.

  • Trends:

    • First ionization energy (IE1) is the energy to remove the first electron from an atom, second ionization energy (IE2) is for removing the second.

    • General expression:

    • For Sodium (Na):

      • Na(g) \rightarrow Na^+(g) + e^-

      • 1^{st} IE = 496 kJ/mol

Trends in First Ionization Energies

  • Generally, less energy is needed to remove the first electron down a group (Z_eff remains constant but distance increases).

  • More energy is required across a period as Z_eff increases.

Exceptions in First Ionization Energy Trends

  • Notable exceptions occur from:

    • 2A to 3A

    • 5A to 6A

    • The case of nitrogen (N) illustrates the increased energy needed due to half-filled sublevels, while oxygen (O) experiences the opposite.

Trends in Successive Ionization Energies

  • Every further electron removed generally costs more energy due to greater attraction of remaining electrons to the nucleus.

  • The energy jumps significantly when starting to remove core electrons compared to valence electrons.

Electron Affinity (EA)

  • Definition: Electron affinity is the energy change for adding an electron to a gaseous atom.

    • Eg.

    • Cl(g) + e^- \rightarrow Cl^-(g)

    • EA = -349 kJ/mol

  • Trends:

    • For alkali metals, electron affinity decreases down the group.

    • Generally, electron affinities become more negative from left to right.

    • Groups 2A and 8A exhibit low EA due to added electrons residing in higher energy orbitals.

Metallic Character

  • Metallic character displays a decreasing trend across periods and increasing trend down groups.

  • It categorizes elements into metals, metalloids, and nonmetals.

Electronegativity

  • Definition: Electronegativity is an atom’s ability to attract electrons within a chemical bond.

    • Fluorine (F) is the most electronegative element.

    • Electronegativity generally increases across a period and decreases down a group.

Summary of Periodic Trends

  • In summary:

    • First ionization energies, electron affinities, and electronegativities generally increase while atomic radii decrease across periods.

    • For groups, trends can vary significantly and should be analyzed in relation to position on the periodic table.