Understanding Electron Configuration and Periodic Table Patterns
What is Electron Configuration?
Electron configuration
Electron configuration refers to the arrangement of electrons in an atom’s energy levels, sublevels, and orbitals.
- Electrons occupy regions around the nucleus called orbitals, which are grouped into sublevels (s, p, d, f) and energy levels (designated by principal quantum numbers, $n$).
- Here’s the key rule: electrons fill orbitals in order of increasing energy, following the Aufbau principle.
- The $1s$ orbital is filled first, followed by $2s$, $2p$, $3s$, and so on.
- Each orbital has a maximum electron capacity: $s$ holds 2, $p$ holds 6, $d$ holds 10, and $f$ holds 14.
Using the Periodic Table to Determine Electron Configuration
The periodic table is divided into blocks(s, p, d, and f) that correspond to the type of sublevel being filled by the outermost electrons of the elements in that block:
- s-block: Groups 1 and 2 (plus helium).
- p-block: Groups 13 to 18.
- d-block: Transition metals (Groups 3 to 12).
- f-block: Lanthanides and actinides.
Determine the electron configuration of phosphorus ($Z=15$).
Solution
- Locate phosphorus in Group 15 and Period 3 of the periodic table.
- Period 3 tells us the outermost electrons are in the $n=3$ energy level.
- Group 15 indicates phosphorus has 5 valence electrons (electrons in the outermost energy level).
- Fill the orbitals in order: $1s^2$, $2s^2$, $2p^6$ (inner electrons sum to 10, corresponding to neon’s configuration).
- The remaining 5 electrons go into $3s^2$ and $3p^3$.
- The full electron configuration is: $1s^2 2s^2 2p^6 3s^2 3p^3$.
- The condensed configuration is: $[Ne] 3s^2 3p^3$.
Use the periodic table as a map: the period number tells you the principal energy level ($n$), while the block and group number guide you to the sublevel and number of valence electrons.
Group Numbers and Valence Electrons
What Do Group Numbers Tell Us?
Group number
The group number of an element indicates the number of valence electrons in its outermost energy level.
Valence electrons are crucial because they determine an element’s chemical reactivity and bonding behavior.
- Group 1 (alkali metals): 1 valence electron ($s^1$).
- Group 2 (alkaline earth metals): 2 valence electrons ($s^2$).
- Groups 13–18 (p-block): The number of valence electrons equals the group number minus 10. For example:
- Group 15 (e.g., nitrogen, phosphorus) has 5 valence electrons ($s^2 p^3$).
- Group 17 (halogens) has 7 valence electrons ($s^2 p^5$).
Consider chlorine ($Z=17$), a Group 17 element:
- Group 17 tells us chlorine has 7 valence electrons.
- Its electron configuration is $[Ne] 3s^2 3p^5$, with 5 electrons in the $3p$ sublevel and 2 in the $3s$ sublevel.
- These 7 valence electrons make chlorine highly reactive, as it only needs 1 more electron to achieve a stable octet.
For transition metals (d-block), the group number corresponds to the total number of electrons in the $s$ and $d$ sublevels of their outermost two energy levels.
Period Numbers and Principal Energy Levels
What Do Period Numbers Tell Us?
Period number
The period number of an element indicates the highest principal energy level ($n$) that contains electrons.
- Period 1: Elements like hydrogen and helium have electrons only in the $n=1$ energy level.
- Period 2: Elements like lithium ($Z=3$) and neon ($Z=10$) have electrons in both the $n=1$ and $n=2$ levels.
- Period 4: Elements like potassium ($Z=19$) and krypton ($Z=36$) have electrons in the $n=1$ to $n=4$ levels.
Let’s analyze bromine ($Z=35$), which is in Period 4 and Group 17:
- Period 4 tells us that bromine’s outermost electrons are in the $n=4$ energy level.
- Group 17 indicates 7 valence electrons.
- Its electron configuration is $[Ar] 4s^2 3d^{10} 4p^5$, with the $4s$ and $4p$ sublevels contributing to the valence electrons.
- Avoid confusing the period number with the number of valence electrons.
- The period number reflects the energy level, not the electron count.
Why Does This Matter?
Understanding electron configurations, group numbers, and period numbers helps explain:
Chemical reactivity: Elements in the same group have similar reactivity because they have the same number of valence electrons.
Alkali metals (Group 1) all react vigorously with water.
Trends across periods: Properties like atomic radius, ionization energy, and electronegativity change predictably across a period because of increasing nuclear charge and constant shielding.
Atomic radius decreases across a period as electrons are pulled closer to the nucleus.
Bonding behavior: The number of valence electrons determines whether an element tends to lose, gain, or share electrons in chemical bonds.
Sodium (1 valence electron) easily loses an electron to form $Na^+$, while chlorine (7 valence electrons) gains an electron to form $Cl^-$
