Atomic Structure, Electronic Configuration, and Ionization Energy

Comprehensive vocabulary flashcards covering atomic structure, electron configurations, subshell rules, periodic and group trends in ionization energy, isotopes, and relative mass terms.

Shell (Energy Level)

The principal quantum energy levels around a nucleus, designated by letters $K$ (1st), $L$ (2nd), $M$ (3rd), and $N$ (4th); the word orbit is never used in this context.

Formula for Maximum Shell Capacity

$2n^2$, where $n$ represents the shell number ($n=1 \rightarrow 2$, $n=2 \rightarrow 8$, $n=3 \rightarrow 18$, $n=4 \rightarrow 32$, $n=5 \rightarrow 50$, $n=6 \rightarrow 72$).

Subshell Electron Capacities

The maximum number of electrons each subshell can hold: $s$ holds $2$, $p$ holds $6$, $d$ holds $10$, and $f$ holds $14$ electrons.

Orbital

The region around the nucleus where the probability of finding an electron is maximum ($95\%$); a single orbital can hold a maximum of $2$ electrons.

Number of Orbitals per Subshell

The $s$ subshell has $1$ orbital, the $p$ subshell has $3$ orbitals, the $d$ subshell has $5$ orbitals, and the $f$ subshell has $7$ orbitals.

Aufbau Principle

The rule stating that electrons will first occupy the subshell or orbital with the lowest energy (nearest to the nucleus).

Pauli Exclusion Principle

The rule stating that an orbital can hold a maximum of $2$ electrons, and they must have opposite spins (one clockwise and one anticlockwise) to cancel out their opposing magnetic fields.

Hund's Rule

The rule stating that when degenerate orbitals of equal energy are available, electrons will fill them singly with parallel (same) spins before pairing up.

Degenerate Orbitals

Orbitals that possess the exact same energy level, such as the three $2p$ orbitals ($2p_x$, $2p_y$, $2p_z$).

Orbital Shapes

An $s$ orbital is spherical, while a $p$ orbital is dumbbell-shaped along three perpendicular axes ($p_x$, $p_y$, $p_z$); the size of both orbital shapes increases as the principal shell number increases.

Subshell Stability Rule

Completely filled subshells ($s^2, p^6, d^{10}, f^{14}$) and half-filled subshells ($s^1, p^3, d^5, f^7$) are stable, whereas partially filled subshells are unstable.

Anomalous Electronic Configuration of Chromium and Copper

Due to the minimal energy gap between $4s$ and $3d$, an electron promotes from $4s$ to $3d$ to achieve stability: Chromium ($Cr$, $Z=24$) forms $[Ar] 4s^1 3d^5$ (half-filled) and Copper ($Cu$, $Z=29$) forms $[Ar] 4s^1 3d^{10}$ (fully filled).

Transition Metal Ionization Order

Although electrons enter the $4s$ subshell before the $3d$ subshell during filling, transition elements ($d$-block) lose their $4s$ electrons first when forming positive ions.

First Ionization Energy ($IE_1$)

The energy required to remove $1\,\text{mole}$ of electrons from $1\,\text{mole}$ of gaseous atoms to form $1\,\text{mole}$ of gaseous positive ions with a $+1$ charge: $\text{Na}(g) \rightarrow \text{Na}^+(g) + 1e^-$.

Second Ionization Energy ($IE_2$)

The energy required to remove $1\,\text{mole}$ of electrons from $1\,\text{mole}$ of gaseous $+1$ ions to form $1\,\text{mole}$ of gaseous ions with a $+2$ charge: $\text{Na}^+(g) \rightarrow \text{Na}^{2+}(g) + 1e^-$.

Endothermic Nature of Ionization Energy

Ionization energy is always an endothermic process because heat energy must be supplied to overcome the electrostatic attractive force holding outer valence electrons to the nucleus.

Requirement for Gaseous Atomic State in IE

The substance must be gaseous to eliminate intermolecular forces (IMF) so energy is solely used to overcome nuclear pull, and atomic because molecular bonds would otherwise require breaking first.

Nuclear Pull (Nuclear Charge)

The positive attraction exerted by the protons in the nucleus; as the number of protons increases, valence electrons are held more tightly, increasing ionization energy ($\text{protons} \propto \text{IE}$).

Shielding Effect (Screening Effect)

The repulsive effect of inner-shell electrons that reduces the effective nuclear attraction felt by outer valence electrons ($\text{shielding electrons} \propto \frac{1}{\text{IE}}$).

Distance Factor in Ionization Energy

The distance between the nucleus and the valence electron; greater distance weakens nuclear attraction, making the electron easier to remove and decreasing ionization energy ($\text{distance} \propto \frac{1}{\text{IE}}$).

Across-a-Period Trend in Atomic Radius

Atomic radius decreases from left to right across a period because the shielding electrons remain constant while the nuclear charge increases, pulling valence electrons closer to the nucleus.

Across-a-Period Trend in Ionization Energy

Ionization energy generally increases across a period because the number of shielding electrons stays constant while the nuclear charge increases and atomic size decreases.

Group 2 to Group 3 Ionization Energy Anomaly

Group 3 elements ($B$, $Al$) have a lower first ionization energy than Group 2 ($Be$, $Mg$) because their outermost electron is in a higher-energy $p$ subshell ($2p$ or $3p$), which is farther from the nucleus and shielded by inner $s$ electrons.

Group 5 to Group 6 Ionization Energy Anomaly

Group 6 elements ($O$, $S$) have a lower first ionization energy than Group 5 ($N$, $P$) due to spin-pair repulsion between two electrons sharing a single $p$ orbital, making one electron easier to remove despite higher nuclear charge.

Down-a-Group Trend in Ionization Energy

First ionization energy decreases down a group because the increasing distance from the nucleus and the added shielding of inner electron shells outweigh the increase in nuclear charge.

Successive Ionization Energies Trend

Successive ionization energies continually increase because fewer electrons are held by the same number of protons; large jumps occur when an electron is removed from a shell closer to the nucleus.

Isotopes

Atoms of the same element containing the same number of protons and electrons but different numbers of neutrons; they share identical chemical properties but differ in physical properties.

Angle of Deflection in an Electric Field

The degree to which a charged particle bends in an electric field, governed by $\text{angle of deflection} \propto \frac{\text{charge}}{\text{mass}}$; electrons deflect significantly more than protons because electrons have a much smaller mass.

Subatomic Particles: Mass, Charge, and Location

Proton has charge $+1$, relative mass $1\,\text{amu}$, located inside nucleus; neutron has charge $0$, relative mass $1\,\text{amu}$, located inside nucleus; electron has charge $-1$, relative mass $\frac{1}{1840} = 0.0005\,\text{amu}$, located outside nucleus.

Relative Atomic Mass ($A_r$)

The weighted average mass of all the isotopes of an element relative to $\frac{1}{12}\text{th}$ of the mass of a carbon-12 atom.

Relative Isotopic Mass

The mass of a single atom of a specific isotope relative to $\frac{1}{12}\text{th}$ of the mass of a carbon-12 atom.

Relative Molecular Mass ($M_r$)

The weighted average of the mass of all atoms in a molecule relative to $\frac{1}{12}\text{th}$ of the mass of a carbon-12 atom; applicable strictly to covalent compounds.

Molecule

A chemical particle formed when two or more atoms are covalently bonded together.

Relative Formula Mass

The weighted average of the mass of a formula unit relative to $\frac{1}{12}\text{th}$ of the mass of a carbon-12 atom; a general term applicable to both ionic and covalent compounds.

Formula Unit

The simplest ratio of oppositely charged ions that make up a giant ionic structure (e.g., $\text{NaCl}$).