Chapter 9: Periodic Properties of the Elements – Electron Configurations and Periodic Trends

CH 9 Periodic Properties of the Elements

Introduction to Periodic Properties

The periodic table organizes elements according to recurring chemical and physical properties, which are explained by their electron configurations. Understanding these trends is essential for predicting element behavior and reactivity.

• Periodic Law: Properties of elements repeat in a predictable way when arranged by increasing atomic number.

• Electron Configuration: The arrangement of electrons in an atom's orbitals determines its chemical properties.

• Valence Electrons: Electrons in the outermost shell are responsible for chemical bonding and reactivity.

Electron Configurations Using the Periodic Table

Electron configurations can be determined systematically using the periodic table and the concept of noble gas shorthand notation.

• Step 1: Locate the element on the periodic table.

• Step 2: Identify the noble gas that precedes the element.

• Step 3: Write the inner electron configuration using the noble gas symbol in brackets (e.g., [Ne]).

• Step 4: Assign electrons to the appropriate orbitals for the elements between the noble gas and the element of interest, following the order of filling across the blocks.

Example: Chlorine (Cl) electron configuration:

• Preceding noble gas: Neon (Ne)

• Configuration: $[Ne] \, 3s^2 3p^5$

Irregular Electron Configurations in Transition Metals

Transition metals (d block) often have electron configurations that deviate from the expected order due to sublevel splitting and energy differences between orbitals.

• Sublevel Splitting: The 4s orbital is lower in energy than the 3d, so it fills first, but the energy difference is small.

• Irregular Configurations: Some transition metals have partially filled ns or (n-1)d orbitals, determined experimentally.

Examples of experimentally determined configurations:

• Cr: $[Ar] \, 4s^1 3d^5$

• Cu: $[Ar] \, 4s^1 3d^{10}$

• Mo: $[Kr] \, 5s^1 4d^5$

• Ru: $[Kr] \, 5s^1 4d^7$

• Pd: $[Kr] \, 5s^0 4d^{10}$

Transition Metal Electron Configurations Table

The electron configurations of transition metals across the d block show systematic filling patterns, with some exceptions.

Element	Electron Configuration
Sc	4s23d1
Ti	4s23d2
V	4s23d3
Cr	4s13d5
Mn	4s23d5
Fe	4s23d6
Co	4s23d7
Ni	4s23d8
Cu	4s13d10
Zn	4s23d10

Additional info: Similar patterns are observed in the 5th period (Mo, Ru, Pd, Ag, Cd).

Periodic Trends and Electron Configuration

Periodic trends in atomic properties are explained by electron configurations and the quantum-mechanical model.

• Elements in the same column have similar chemical properties due to similar valence electron configurations.

• Elements in a period show repeating patterns in properties.

• Reactivity is highest for elements with electron configurations closest to noble gases.

• Atoms tend to gain or lose electrons to achieve noble gas configurations.

Noble Gas Electron Configuration

Noble gases have full valence shells, making them especially stable and chemically inert.

• Most noble gases have eight valence electrons (except He, which has two).

• Other elements react to achieve this stable configuration.

Alkali Metals and Halogens: Electron Configurations and Reactivity

Alkali metals and halogens are highly reactive due to their proximity to noble gas configurations.

• Alkali Metals (Group 1): One electron more than the previous noble gas; tend to lose one electron to form $1^+$ cations.

• Halogens (Group 17): One electron fewer than the next noble gas; tend to gain one electron to form $1^-$ anions.

Ion Formation and Electron Configuration

Many metals and nonmetals form ions with predictable charges based on their group position.

• Group 1A: $1^+$

• Group 2A: $2^+$

• Group 7A: $1^-$

• Group 6A: $2^-$

Ion formation results in electron configurations matching the nearest noble gas.

Atomic Radius

Atomic radius is a measure of the size of an atom, which varies predictably across the periodic table.

• Van der Waals Radius: Half the distance between nuclei of identical atoms not bonded together.

• Covalent Radius: Half the distance between nuclei of identical atoms bonded together.

• Trends:

  • Decreases across a period (left to right) due to increasing nuclear charge pulling electrons closer.

  • Increases down a group due to addition of electron shells.

Trends in Atomic Radius: Transition Metals

Transition metals show less variation in atomic radius across a period compared to main-group elements.

• Atomic radii of transition metals are roughly constant across the d block.

• Electron number remains fairly constant, so size does not change much.

Ionic Radius

Ionic radius refers to the size of an ion, which depends on its charge and electron configuration.

• Cations are smaller than their neutral atoms; anions are larger.

• For isoelectronic species (same number of electrons), higher positive charge means smaller cation, higher negative charge means larger anion.

Ionization Energy (IE)

Ionization energy is the energy required to remove an electron from an atom or ion in the gaseous state.

• First ionization energy ($IE_1$): $M(g) + E_1 \rightarrow M^+(g) + e^-$

• Second ionization energy ($IE_2$): $M^+(g) + E_2 \rightarrow M^{2+}(g) + e^-$

• IE increases across a period and decreases down a group.

• Large jump in IE when removing core electrons after valence electrons are gone.

Electron Affinity (EA)

Electron affinity is the energy change when an atom gains an electron in the gaseous state.

• Defined as exothermic: $M(g) + e^- \rightarrow M^-(g) + EA$

• More negative EA means the atom more readily accepts an electron.

• EA generally increases (becomes more negative) across a period; decreases down a group.

• Halogens have the highest (most negative) electron affinities.

Metallic and Nonmetallic Character

Elements are classified as metals, nonmetals, or metalloids based on their physical and chemical properties.

Property	Metals	Nonmetals
Appearance	Lustrous, reflective	Dull, nonreflective
Conductivity	Good conductors	Poor conductors
Malleability/Ductility	Malleable, ductile	Brittle
Oxide Behavior	Basic oxides	Acidic oxides
Ion Formation	Form cations	Form anions

• Metallic character increases down a group and decreases across a period.

Summary of Periodic Trends

Periodic trends are summarized as follows:

• Atomic radius: Increases down a group, decreases across a period.

• Ionization energy: Decreases down a group, increases across a period.

• Electron affinity: Decreases down a group, increases across a period.

• Metallic character: Increases down a group, decreases across a period.

Additional info: These trends are explained by changes in nuclear charge, electron shielding, and quantum-mechanical principles.