Ionization Energy
Ionization energy (IE) is the minimum energy required to remove the most loosely bound electron from a gaseous atom or ion in its ground state, producing a positive ion. The first ionization energy (IE₁) removes the first electron; successive ionization energies increase because each removal leaves behind a more positively charged species that holds remaining electrons more tightly. Ionization energy increases across a period (due to greater effective nuclear charge) and decreases down a group (due to greater atomic radius and electron shielding), making it a key periodic trend.
Key Formula
X(g) → X⁺(g) + e⁻ ΔH = IE₁
LaTeX: X(g) \rightarrow X^+(g) + e^- \quad \Delta H = IE_1
| Symbol | Meaning | Unit |
|---|---|---|
| X(g) | Neutral gaseous atom | — |
| X⁺(g) | Singly charged gaseous cation | — |
| e⁻ | Removed electron | — |
| IE₁ | First ionization energy | kJ mol⁻¹ or eV |
Worked Example
Problem
The first ionization energy of sodium is 496 kJ mol⁻¹. How much energy in kJ is needed to ionize 2.5 mol of gaseous sodium atoms?
Solution
Step 1 — Identify known values: IE₁(Na) = 496 kJ mol⁻¹ n(Na) = 2.5 mol Step 2 — Multiply IE₁ by the number of moles: Energy = IE₁ × n = 496 kJ mol⁻¹ × 2.5 mol Step 3 — Calculate: Energy = 1240 kJ
Answer
1240 kJ of energy is required.
First Ionization Energies of Period 3 Elements (kJ mol⁻¹)
| Element | Symbol | Z | IE₁ (kJ mol⁻¹) | Trend |
|---|---|---|---|---|
| Sodium | Na | 11 | 496 | Lowest in period |
| Magnesium | Mg | 12 | 738 | Increasing → |
| Aluminium | Al | 13 | 578 | Dip (3p vs 3s) |
| Silicon | Si | 14 | 786 | Increasing → |
| Phosphorus | P | 15 | 1012 | Increasing → |
| Chlorine | Cl | 17 | 1251 | Near highest in period |
Interactive Tools
Ptable Ionization Energy Trend
Interactive periodic table displaying ionization energies for all elements.
Open Tool
Wikimedia Commons, CC BY-SA
Related Terms
Electron Affinity
Electron affinity (EA) is the energy change that occurs when a neutral gaseous atom gains one electron to form a negative ion (anion). A negative EA value indicates an exothermic process — the atom releases energy and the anion is more stable than the separated atom and electron — which is the case for most halogens. Electron affinity generally increases (becomes more negative) across a period and decreases down a group, though there are notable exceptions such as the anomalously low EA of fluorine compared to chlorine due to electron–electron repulsion in fluorine's compact 2p orbitals.
Effective Nuclear Charge
Effective nuclear charge (Z_eff) is the net positive charge experienced by a valence electron after accounting for the shielding (screening) effect of inner electrons, which partially cancel the attraction from the nucleus. It is calculated as Z_eff = Z − S, where Z is the actual atomic number and S is the shielding constant. Effective nuclear charge increases across a period because additional protons are added while shielding remains approximately constant, explaining trends in atomic radius, ionization energy, and electron affinity.
Atomic Radius
Atomic radius is a measure of the size of an atom, typically defined as half the distance between the nuclei of two identical adjacent atoms in a covalent bond (covalent radius) or in a metallic lattice (metallic radius). Atomic radius decreases across a period (left to right) because increasing nuclear charge pulls electrons closer to the nucleus, while it increases down a group because additional electron shells increase the average distance of the outermost electrons from the nucleus. These periodic trends directly influence bond lengths, ionic sizes, and many physical properties.
From Greek "ion" (going, from "ienai" to go) referring to charged particles, and Latin "energia" (activity, operation). The concept was developed alongside Arrhenius's theory of electrolytic dissociation in the 1880s; ionization energies were measured spectroscopically in the early 20th century.