Source: General Chemistry lecture notes (Purdue University) | Topic: Periodic Trends, Ions, Ionic Bonding
Tags: atomic radius, periodic trends, cation, anion, isoelectronic, ionic compounds, ion size, noble gas configuration, metallic bonding, covalent bonding
Difficulty: Introductory Prerequisites: Basic understanding of atomic structure (protons, neutrons, electrons), electron configuration notation, and familiarity with the layout of the periodic table.
Big Picture
This topic covers how and why atoms differ in size across the periodic table, and what happens to that size when atoms gain or lose electrons to become ions. It sits at the intersection of atomic structure and chemical bonding, which means it feeds directly into understanding why compounds form the way they do. If you are comfortable with electron configurations and know your way around the periodic table's groups and periods, you are ready for this. If not, revise those first, as every trend here depends on where an element sits on the table.
Atoms get larger as you move down a group (more electron shells) and smaller as you move left to right across a period (stronger nuclear pull on the same shell). When atoms lose electrons they become cations and shrink; when they gain electrons they become anions and grow. Ions with the same electron count are called isoelectronic, and comparing their sizes comes down to how many protons each nucleus has.
Atomic radius
The distance from the centre of an atom's nucleus to the outer edge of its electron cloud. Measured in picometres (1 pm = 10⁻¹² m). In simple terms, this is the "size" of an atom.
Cation
A positively charged ion, formed when a neutral atom loses one or more electrons. Think of it as an atom that has given away electrons and now has more protons than electrons.
Anion
A negatively charged ion, formed when a neutral atom gains one or more electrons. Think of it as an atom that has picked up extra electrons and now carries a negative charge.
Isoelectronic
Describes a group of ions (or atoms) that share the same total number of electrons. In simple terms, isoelectronic species have identical electron configurations even though they are different elements.
Valence electrons
The electrons in the outermost energy level (n) of an atom. These are the electrons involved in bonding and the ones lost first when a cation forms.
Noble gas configuration
The stable electron arrangement that matches one of the noble gases (He, Ne, Ar, Kr, Xe, Rn). Metals lose electrons and non-metals gain electrons to reach the nearest noble gas configuration.
Ionic compound
A compound formed when a metal transfers electrons to a non-metal, creating cations and anions held together by electrostatic attraction. Think of it as a structure built from oppositely charged ions locked in a crystal lattice.
Metallic bond
The bond formed when two metals combine, sharing a "sea" of delocalised electrons.
Covalent bond
The bond formed when two non-metals share electrons rather than transferring them.
Atomic radius increases as you move down a group.
Each new period adds an electron shell, so the outermost electrons sit further from the nucleus.
Example: Li (152 pm) → Na (186 pm) → K (227 pm).
Atomic radius decreases as you move left to right across a period.
Protons are added to the nucleus with each element, increasing the effective nuclear charge.
The extra pull draws the electron cloud inward, even though electrons are being added to the same shell.
Example: Na (186 pm) → Mg (160 pm) → Al (143 pm) → Cl (99 pm).
Bottom-left of the periodic table holds the largest atoms (e.g. caesium, francium).
Top-right holds the smallest atoms (e.g. fluorine, helium is smallest overall but noble gases are often excluded from radius comparisons).
A neutral atom has equal numbers of protons and electrons.
Cation formation (metals)
A metal atom loses one or more valence electrons.
It loses as many as needed to reach the electron configuration of the nearest previous noble gas.
Example: Na (1s² 2s² 2p⁶ 3s¹) loses 1 electron → Na⁺ (1s² 2s² 2p⁶), matching neon.
Electrons are always lost from the highest principal energy level (n) first.
Anion formation (non-metals)
A non-metal atom gains electrons to reach the electron configuration of the next noble gas.
Example: Cl (1s² 2s² 2p⁶ 3s² 3p⁵) gains 1 electron → Cl⁻ (1s² 2s² 2p⁶ 3s² 3p⁶), matching argon.
Cations are smaller than their parent atoms.
Losing electrons often removes an entire outer shell, and the remaining electrons are pulled closer by the same nuclear charge.
Anions are larger than their parent atoms.
Gaining electrons increases electron-electron repulsion in the outer shell, causing the electron cloud to expand.
The group trend (larger as you go down) still holds for ions.
The period trend for neutral atoms does not carry over neatly to ions, because cations and anions belong to different isoelectronic series.
Isoelectronic ions all have the same number of electrons.
Among isoelectronic species, the ion with the most protons is the smallest (greater nuclear charge pulls the shared electron cloud tighter).
Example (all 10 electrons, neon configuration): N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺. Nitrogen has fewest protons (7), aluminium has most (13), so N³⁻ is largest and Al³⁺ is smallest.
Metal + metal → metallic bond
Metal + non-metal → ionic bond
Non-metal + non-metal → covalent bond
Ionic compounds form crystalline solids that are:
Hard and brittle
High melting and boiling points
Good electrical conductors when molten or dissolved in solution (ions are free to move)
Generally soluble in water
"Atomic size increases across a period from left to right." It does the opposite. More protons pull the electron cloud inward, so atoms get smaller from left to right. The increase happens going down a group.
"Cations are bigger than their neutral atoms because they are charged." Cations are smaller. Losing electrons (often an entire shell) means the remaining electrons sit closer to the nucleus.
"Electrons are lost from the inner shells first." Electrons are always removed from the highest principal energy level (the valence shell) first, even for transition metals.
"All ions of the same element are the same size." The charge matters. Fe²⁺ and Fe³⁺ are different sizes because Fe³⁺ has lost one more electron, giving it a smaller radius.
"Isoelectronic ions are all the same size." They have the same number of electrons, but the ion with more protons will be smaller. Nuclear charge is what differentiates them.
⚠️ Ranking atoms or ions by size is one of the most common periodic-trends exam questions. Know both directions: down a group (bigger) and across a period (smaller).
⚠️ Isoelectronic comparisons come up frequently. The method is simple: same electron count, more protons = smaller ion. Be ready to rank a set like O²⁻, F⁻, Na⁺, Mg²⁺.
⚠️ You will likely be asked which electrons are removed first when forming a cation. The answer is always the valence (highest n) electrons.
⚠️ Know the three bond-type rules cold: metal + metal = metallic, metal + non-metal = ionic, non-metal + non-metal = covalent. These appear as quick identification questions.
⚠️ Properties of ionic compounds (high melting point, brittle, conduct when molten/dissolved) are a staple of short-answer and multiple-choice sections.
Ionic size trends matter in materials science and engineering. The relative sizes of cations and anions determine how ions pack into crystal lattices, which in turn controls the mechanical strength, melting point, and solubility of salts. Table salt (NaCl) dissolves easily in water partly because the Na⁺ and Cl⁻ ions are small enough for water molecules to pull apart and surround individually.
True or false: Atomic radius increases from left to right across a period.
Answer: False. It decreases because of increasing nuclear charge.
Fill in the blank: A positively charged ion is called a ______.
Answer: Cation.
True or false: An anion is larger than its parent neutral atom.
Answer: True. The extra electron(s) increase repulsion and expand the electron cloud.
Fill in the blank: Among isoelectronic species, the ion with the most ______ is the smallest.
Answer: Protons.
True or false: When sodium forms Na⁺, it loses an electron from the 1s orbital.
Answer: False. It loses the electron from the 3s (valence) level.
Q: Arrange the following in order of increasing atomic radius: Na, Li, K.
A: Li < Na < K. All are in Group 1; atomic radius increases going down the group as more electron shells are added.
Q: Which is larger, a chlorine atom or a chloride ion (Cl⁻)? Explain why.
A: Cl⁻ is larger. Gaining an extra electron increases electron-electron repulsion in the outer shell, expanding the electron cloud, while the nuclear charge stays the same.
Q: Rank the following isoelectronic species from smallest to largest: O²⁻, F⁻, Na⁺, Mg²⁺.
A: Mg²⁺ < Na⁺ < F⁻ < O²⁻. All have 10 electrons. More protons means a tighter pull on those electrons, so Mg²⁺ (12 protons) is smallest and O²⁻ (8 protons) is largest.
Q: What type of bond forms between magnesium and chlorine? What properties would you expect the resulting compound to have?
A: Ionic bond (metal + non-metal). MgCl₂ would be a crystalline solid, hard and brittle, with a high melting point, and it would conduct electricity when molten or dissolved in water.
Q: A student claims that when iron forms Fe²⁺, it first loses two electrons from the 3d subshell. Is this correct?
A: No. Electrons are removed from the highest principal energy level first. Iron loses its two 4s electrons before any 3d electrons, giving Fe²⁺ the configuration [Ar] 3d⁶.
This material connects directly to electron configuration (you need it to predict how many electrons an atom will lose or gain) and to ionisation energy and electronegativity, which are the other major periodic trends tested alongside atomic radius. Understanding ion sizes also feeds into lattice energy calculations later in the course, where the distance between ions in a crystal determines the strength of the ionic bond.
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