lecture 11

Electron Configuration, Valence Electrons, and the Periodic Table

Core and Valence Electrons

Understanding the arrangement of electrons in atoms is fundamental to predicting chemical behavior. Electrons are organized into shells and subshells, with those in the outermost shell playing a key role in chemical reactions.

• Core (inner) electrons: Electrons in lower-energy shells, not involved in bonding.

• Valence electrons: Electrons in the outermost shell (highest principal energy level, n), responsible for chemical reactivity and bonding.

• The number of valence electrons determines an element's chemical and physical properties.

• Valence electrons are lost to form cations or gained/shared to form anions or covalent bonds.

Example: For silicon (Si), the electron configuration is $1s^2 2s^2 2p^6 3s^2 3p^2$. The 3s and 3p electrons (total of 4) are valence electrons; the rest are core electrons.

Determining Core and Valence Electrons: Worked Example

To determine the number of core and valence electrons, write the full electron configuration and identify electrons in the highest principal energy level as valence electrons.

• Example (Germanium, Ge): Atomic number 32. Electron configuration: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^2$

• Valence electrons: Electrons in $n=4$ shell ($4s^2 4p^2$) = 4 valence electrons.

• Core electrons: All others ($1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10}$) = 28 core electrons.

Practice: Write the electron configuration for phosphorus (P) and identify its valence and core electrons.

Orbital Blocks and Their Position in the Periodic Table

Periodic Table Organization

The periodic table is divided into four blocks based on the type of atomic orbital being filled:

• s-block: Groups 1 and 2, plus helium

• p-block: Groups 13–18

• d-block: Transition metals (Groups 3–12)

• f-block: Lanthanides and actinides

• The group number of a main-group element equals its number of valence electrons.

• The row number (period) equals the highest principal quantum number (n) for that element.

Using the Periodic Table to Write Electron Configurations

Electron configurations can be written using the periodic table as a guide:

1. Locate the element on the periodic table.

2. Find the noble gas that precedes the element; write its symbol in brackets to represent core electrons.

3. Continue by adding electrons to the appropriate orbitals as you move across the table to the element of interest.

Example (Selenium, Se): Atomic number 34. Preceding noble gas is argon [Ar]. Configuration: [Ar] $4s^2 3d^{10} 4p^4$

Transition and Inner Transition Metals

Electron Configurations of Transition Metals

Transition metals (d-block) and inner transition metals (f-block) have unique electron configurations due to sublevel energy differences.

• The $4s$ orbital fills before the $3d$ orbital, but the energy difference is small.

• Some transition metals have irregular configurations (anomalies) to achieve half-filled or fully filled d sublevels.

Examples:

• Chromium (Cr): [Ar] $4s^1 3d^5$ (instead of $4s^2 3d^4$)

• Copper (Cu): [Ar] $4s^1 3d^{10}$ (instead of $4s^2 3d^9$)

Table: Selected Transition Metal Electron Configurations

Element	Expected Configuration	Actual Configuration
Cr	[Ar] $4s^2 3d^4$	[Ar] $4s^1 3d^5$
Cu	[Ar] $4s^2 3d^9$	[Ar] $4s^1 3d^{10}$
Mo	[Kr] $5s^2 4d^4$	[Kr] $5s^1 4d^5$
Ag	[Kr] $5s^2 4d^9$	[Kr] $5s^1 4d^{10}$

Electron Configuration and Elemental Properties

Periodic Patterns

Elements in the same group (vertical column) have similar properties due to the same number of valence electrons. Properties repeat periodically across periods (rows).

Noble Gases

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

• They are chemically inert due to their stable, full valence shells.

Metals

• Make up the majority of elements.

• Alkali metals (Group 1): Lose one electron to form $1^+$ cations.

• Alkaline earth metals (Group 2): Lose two electrons to form $2^+$ cations.

• Transition and inner transition metals: Lose electrons from s and then d orbitals; can form multiple cations.

• p-block metals: Lose electrons from s and p orbitals to form cations.

Metalloids

• Located along the zigzag line between metals and nonmetals in the p-block.

• Can exhibit both metallic and nonmetallic behavior.

• May lose electrons to form cations or gain electrons to form anions.

Nonmetals

• Located in the upper right of the periodic table (mainly p-block).

• Gain electrons in p orbitals to form anions with noble gas configurations.

Halogens

• Nonmetals with one fewer electron than the next noble gas.

• Gain one electron to form $1^-$ anions in reactions with metals.

• Share electrons with other nonmetals to achieve noble gas configurations.

Ion Formation and Predictable Charges

Predicting Ion Charges

The charge of ions formed by main-group elements can be predicted from their group number and electron configuration.

• Atoms form ions to achieve the electron configuration of the nearest noble gas.

• Metals form cations (positive ions); nonmetals form anions (negative ions).

• Alkali metals (Group 1A): $1^+$ cations

• Alkaline earth metals (Group 2A): $2^+$ cations

• Halogens (Group 7A): $1^-$ anions

Example: Aluminum (Al) loses three electrons to form $Al^{3+}$; sulfur (S) gains two electrons to form $S^{2-}$.

Table: Predictable Ion Charges for Main-Group Elements

Group	Typical Ion Charge
1A	+1
2A	+2
3A	+3
5A	-3
6A	-2
7A	-1

Periodic Trends: Atomic Radii and Effective Nuclear Charge

Atomic Radius

• Increases down a group (valence shell farther from nucleus).

• Decreases across a period (left to right) as electrons are added to the same shell and the nucleus pulls them closer.

Effective Nuclear Charge ($Z_{eff}$)

The effective nuclear charge is the net positive charge experienced by valence electrons, accounting for shielding by core electrons.

• Core electrons shield outer electrons from the full nuclear charge.

• Valence electrons do not shield each other effectively.

• $Z_{eff}$ increases across a period and decreases down a group.

• Electrons in s orbitals shield more effectively than those in p orbitals.

Formula: $Z_{eff} = Z - S$ Where $Z$ is the atomic number and $S$ is the number of core electrons (shielding constant).

Summary Table: Periodic Trends

Trend	Down a Group	Across a Period
Atomic Radius	Increases	Decreases
Effective Nuclear Charge	Decreases	Increases

Additional info: Some context and explanations have been expanded for clarity and completeness, including the summary tables and explicit formulas.