Periodic Properties of the Elements

Periodic Properties of the Elements

Module Learning Goals

• By the end of this module, students will be able to:

  1. Write electron configurations from the periodic table and relate quantum numbers to the location of elements in the periodic table.

  2. Estimate the effective nuclear charge, $Z_{eff}$, and use it to explain and predict trends in:

  • i. Atomic size

  • ii. Ionic size

  • iii. Relative ionization energies

  • iv. Electron affinity

  1. Recognize periodic behavior of the elements.

  2. Identify the three main types of chemical bonds and how to classify them based on electronegativity.

Development of the Periodic Table

• Early Classifications:

  • Ancient Times: Knowledge of elements was rudimentary.

  • Middle Ages (1700): Development of more organized classifications.

  • 1735-1843: Various attempts to group elements.

  • 1843-1886: Identification of new elements; atomic mass became a primary categorization criterion.

  • 1894-1918: Further classification and recognition of periodic trends.

  • 1923-1961: The periodic table was refined to better reflect periodic properties; significant developments and discoveries occurred.

  • 1965-Present: Continuous updates and new elements added.

• John Newlands (1865):

  • Arranged elements in order of atomic mass and noted that every eighth element had similar properties, termed
    "Law of Octaves."

  • This trend was effective only up to calcium.

• Dmitri Mendeleev (1870):

  • Arranged elements by atomic mass and made adjustments to group elements with similar properties.

  • Left "gaps" in the table to accommodate undiscovered elements.

  • Predicted properties of elements that were not yet discovered by assessing trends in known elements.

Mendeleev's Periodic Table

• Notable Features:

  • Predicted the existence of an element below aluminum based on properties.

  • Example prediction included the existence of germanium.

Quantum Mechanics and Periodicity

• Before 1927, the reasons for periodic trends were unknown.

• Quantum Mechanics validates Mendeleev’s observations.

• Ordering the periodic table by atomic number $Z$ instead of atomic mass revealed similar periodic properties, confirming Mendeleev’s initial findings.

• Henry Moseley (1913): Developed the concept that atomic number is the proper organizing principle for the periodic table, leading to predictions consistent with Mendeleev's.

Alkali Metals and Chemical Properties

• Alkali Metals Properties:

  • Characteristics: Soft metals, low melting points, reactive with water.

  • Reactivity pattern: Increases as you move down the group.

  • Reaction Equation: $2 M(s) + 2 H2O(l) ightarrow 2 MOH(aq) + H2(g)$.

• Periodic Trends in Reactivity with Oxygen:

  • How elements across a period react with O$_2$ shows observable patterns influenced by electron distribution in orbitals.

Multi-Electron Systems and Orbitals

• Orbital Energies: Depend on principal quantum number $n$ and angular momentum $l$.

  • Example: The 4s orbital is lower in energy than the 3d orbital.

Origin of Periodic Trends: Electron Configuration

• Valence Electrons: Outermost electrons primarily involved in chemical bonding.

• Core Electrons: Inner electrons corresponding to previous noble gases.

Metallic, Non-Metallic, and Semi-Metallic Classification

• Metallic Character: Refers to how closely an element's properties resemble ideal metals (e.g., being malleable, ductile, conductive).

• Trends in Metallic Character:

  • Decreases left to right across periods.

  • Increases down groups.

• Metals Tend to Lose Electrons:

  • Example Reactions:

  • $2 K(s) + 2 H2O(l) ightarrow 2 K^+(aq) + 2 OH^-(aq) + H2(g)$

  • $2 Na(s) + 2 H2O(l) ightarrow 2 Na^+(aq) + 2 OH^-(aq) + H2(g)$

• Common Electron Configurations upon Ionization:

  • $Na: [Ne] 3s^1
    ightarrow Na^+: [Ne]$

  • $K: [Ar] 4s^1
    ightarrow K^+: [Ar]$

Nonmetals Tend to Gain Electrons

• Example Configurations:

  • Chlorine: $Cl: [Ne] 3s^2 3p^5
    ightarrow Cl^-: [Ne] 3s^2 3p^6$.

• Typical Reactions:

  • $2 Al(s) + 3 Cl2(g) ightarrow 2 AlCl3(s)$

  • $2 Al(s) + 3 Br2(g) ightarrow 2 AlBr3(g)$

Trends in Ionic Radius

• Ionic Radius:

  • Cations are smaller than neutral atoms due to loss of valence electrons.

  • Example: $S^{2-} > Cl^- > K^+ > Ca^{2+}$

  • Anions are larger than neutral atoms due to gain of electrons.

Ionization Energy (IE)

• Ionization Energy Definition: Minimum energy required to remove an electron from an atom or ion, endothermic process.

• Common Equations for Ionization:

  • $M(g) + IE_1
    ightarrow M^+(g) + e^-$

  • $M^+(g) + IE_2
    ightarrow M^{2+}(g) + e^-$

• Trends in Ionization Energy:

  • Decreases down a group, increases across a period due to increased effective nuclear charge (Z) acting on valence electrons.

Electron Affinity (EA)

• Electron Affinity Definition: Energy released when an atom gains an electron, can either be positive or negative.

• Generally: EA becomes less negative down a column and more negative across a row.

Chemical Reactivity and Bond Formation

• Reactive Atoms: Atoms with incomplete valence shells try to achieve full shells by gaining, losing, or sharing electrons.

Types of Chemical Bonds

• Ionic Bonds: Form between metals and nonmetals (large electronegativity difference), atoms do not share electrons.

• Covalent Bonds: Form between nonmetals where electrons are shared; can be polar or non-polar depending on electronegativity differences.

Electronegativity Trends

• General Trend: Electronegativity increases from left to right across a period and decreases down a group.

• Pauling Electronegativity Values: A numerical scale developed by Linus Pauling to quantify electronegativity differences.

Summary

• The periodic table organizes elements based on increasing atomic number, showing trends in atomic and ionic size, ionization energy, electron affinity, and bond types.

• General categories include metals, nonmetals, and semi-metals based on elemental properties and behaviors.