Periodic Properties and Trends

Note on topic scope: This topic (T13) covers periodic properties and trends — using the periodic table to predict atomic and ionic size, ionization energy, electronegativity, electron affinity, and bond-type tendencies. If you need to write or interpret electron configurations, start with T12: Electron Structure and Light; come back here to reason about how those configurations drive periodic patterns.

Valence electrons are the electrons in the outermost energy level (highest n value) of an atom. They are the electrons that participate in bonding and chemical reactions, and they determine an element’s chemistry.

• For main-group elements, the number of valence electrons equals the column position within the s/p block. Group 1 has 1 valence electron (ns⁠¹); Group 2 has 2 (ns⁠²); Group 13 has 3 (ns⁠² np⁠¹); … Group 17 has 7 (ns⁠² np⁠⁵); Group 18 has 8 (ns⁠² np⁠⁶), except He which has 2.
• For transition metals, both the ns and (n−1)d electrons can participate in chemistry, which is why transition metals show variable oxidation states.

The periodic table is organized by electron configuration, which is what makes it predictive:

• s-block — Groups 1 and 2 (ns⁠¹ and ns⁠² configurations).
• p-block — Groups 13–18 (ns⁠² np⁠¹ … ns⁠² np⁠⁶ configurations).
• d-block — transition metals (filling the (n−1)d subshell).
• f-block — lanthanides and actinides (filling the (n−2)f subshell).

The period number equals the highest principal quantum number of the valence electrons. This direct connection between table position and configuration means you can write any element’s configuration just by reading its position.

Every periodic trend comes back to a single tug-of-war: nuclear pull on valence electrons versus shielding by inner-shell electrons. The net pull on a specific electron is its effective nuclear charge, Z_eff: the pull exerted on that electron by the nucleus, taking electron-electron repulsions into account. A useful approximation is Z_eff ≈ Z − (number of inner-shell electrons). Core electrons are adept at shielding, while electrons in the same valence shell do not block the nuclear attraction experienced by each other as efficiently.

• Across a period (left to right): Z increases by 1 for each step, but the shielding from inner-shell electrons stays roughly constant. So Z_eff increases across a period — the nucleus pulls valence electrons more tightly.
• Down a group (top to bottom): each new period adds a full inner shell of electrons. The valence electrons sit in a higher shell, farther from the nucleus, with much more shielding. So Z_eff changes only modestly down a group, but the distance to the valence shell grows substantially.

Three of the four periodic trends — atomic radius, ionization energy, electronegativity — follow directly from these two patterns. Memorizing the arrows is fine; the better skill is to reason through Z_eff + distance and let the answer fall out.

Atomic radius measures the size of an atom — typically the distance from the nucleus to the outermost electron shell.

• Across a period (left to right): radius decreases. Z_eff increases (more protons pulling), shielding stays roughly constant (same shell), so the valence electrons are pulled in closer.
• Down a group (top to bottom): radius increases. Each new period adds a shell, which puts the outermost electrons farther from the nucleus.

Examples:

• Across period 3: Na > Mg > Al > Si > P > S > Cl (decreasing).
• Down group 1: Li < Na < K < Rb < Cs (increasing).

To rank a mixed set of elements (e.g., Na vs Cl vs Cs), apply both rules: down-and-left is bigger; up-and-right is smaller. Cs is large (heavy alkali metal); Cl is small (right side, period 3); Na sits between them.

Ionization energy (IE) is the energy required to remove the most loosely bound electron from a gaseous atom in its ground state: X(g) → X⁠⁺(g) + e⁠^-. The first ionization energy (IE₁) removes the first electron; IE₂, IE₃, … remove subsequent electrons. Ionization processes are endothermic, so IE values are always positive.

The general trend mirrors atomic radius (because closer-held electrons are harder to remove):

• Across a period: IE increases. Higher Z_eff, smaller radius → harder to remove the valence electron.
• Down a group: IE decreases. Larger radius and more shielding → the outermost electron is loosely held.

Major exceptions. Two well-known dips break the smooth left-to-right increase across a period:

1. Group 13 vs Group 2 (e.g., B < Be, Al < Mg). Group 13 starts the np subshell. Within any one shell, the s electrons are lower in energy than the p electrons, so an s electron is harder to remove than a p electron in the same shell. The first p electron is also shielded by the filled ns⁠², so it leaves more easily than expected.
2. Group 16 vs Group 15 (e.g., O < N, S < P). Group 15 has a half-filled np⁠³ — an extra-stable arrangement. Removing one electron from the group-16 element eliminates the electron-electron repulsion caused by pairing electrons in a 2p orbital and leaves a half-filled p subshell (which is energetically favorable), so this electron leaves more easily than the smooth trend predicts.

For successive ionization energies (IE₁, IE₂, IE₃, …) of a single element, look for the large jump when you start removing core electrons. Sc and Ga both have 3 valence electrons; both show a sharp increase in IE between IE₃ and IE₄ as the removal hits the noble-gas core. The position of the big jump tells you how many valence electrons the atom has.

Two more energy-related trends share the same direction as ionization energy:

Electronegativity (EN) is a relative scale of an atom’s tendency to attract shared electrons in a bond. The Pauling scale is the most common; values increase across a period and decrease down a group, for the same Z_eff-and-distance reasons that govern atomic radius and ionization energy. Fluorine has the highest electronegativity on the Pauling scale; the heavy alkali metals (Cs, Fr) are the lowest. EN is most useful as a way to predict bond polarity and bond-type tendency (see the LO 13.6 section below).

Electron affinity (EA) is the energy change when a gaseous atom gains an electron to form an anion: X(g) + e⁠^- → X⁠^-(g). Conventionally a negative EA value means energy is released (exothermic), so “more negative EA” means “more eager to accept an electron.”

• Across a period: EAs tend to become more negative (more exothermic) as Z_eff increases.
• Down a group: EAs tend to become less negative.
• Halogens have the most negative EA values (chlorine’s EA is −348 kJ/mol, the most negative of any element).
• Noble gases (group 18), group 2 elements, and group 15 (because of half-filled np⁠³ stability) show exceptions to the trend — for these groups, adding an electron is much less favorable than the smooth left-to-right pattern would predict, because the incoming electron has to occupy a higher-energy subshell or pair against an existing half-filled p orbital.
• Counter-intuitively, fluorine’s EA (−328 kJ/mol) is less negative than chlorine’s, even though F is the most electronegative element. The reason is the very small n = 2 valence shell of fluorine: the incoming electron experiences strong electron-electron repulsion in the cramped 2p shell, which the larger 3p shell of chlorine avoids.

The same Z_eff + distance reasoning that explained IE governs EN and EA as well. Think of the three (IE, EN, EA-magnitude) as one “hold on to electrons” trend with three different operational definitions.

Forming an ion changes the size dramatically:

• Cations are smaller than their parent atoms. A cation has fewer electrons and the same number of protons as the parent atom, so it is smaller than the atom from which it is derived. As electrons are removed from the outer valence shell, the remaining core electrons (now occupying smaller shells) experience a greater effective nuclear charge Z_eff and are drawn even closer to the nucleus. Often a whole shell is lost: Al has a covalent radius of 118 pm; Al⁠³⁺ has an ionic radius of 68 pm.
• Anions are larger than their parent atoms. Adding electrons increases electron-electron repulsion and spreads the cloud out: a sulfur atom has a covalent radius of 104 pm, while the sulfide anion has an ionic radius of 170 pm.
• For cations of the same element with different charges, the higher charge gives the smaller ion: V⁠²⁺ has an ionic radius of 79 pm versus V⁠³⁺ at 64 pm.

Isoelectronic species are atoms and ions that have the same number of electrons (and therefore the same electron configuration). Examples include N⁠³−, O⁠²−, F−, Ne, Na⁠⁺, Mg⁠²⁺, and Al⁠³⁺, which all have the Ne configuration, 1s⁠²2s⁠²2p⁠⁶. The size rule for an isoelectronic series is simple: more protons = smaller radius. For the 10-electron series:

N⁠³− (Z=7) > O⁠²− (Z=8) > F− (Z=9) > Ne (Z=10) > Na⁠⁺ (Z=11) > Mg⁠²⁺ (Z=12) > Al⁠³⁺ (Z=13)

All seven species have the same [Ne] electron arrangement, but Al⁠³⁺ has 13 protons pulling on those 10 electrons, while N⁠³− has only 7. The size ranking is purely a Z_eff argument with everything else held constant.

Same-period cation series vs same-period anion series. When the metals on the left of a period (e.g., Na, Mg, Al) form their typical cations, they lose their entire valence shell, so Na⁠⁺, Mg⁠²⁺, Al⁠³⁺ all share the [Ne] configuration of period 2. The nonmetals on the right of the same period (P, S, Cl) form anions by completing their valence shell, so P⁠³−, S⁠²−, Cl− all share the [Ar] configuration of period 3. The result: every typical anion of a period sits a full shell larger than every typical cation of the same period. P⁠³−, S⁠²−, Cl− are all larger than Na⁠⁺, Mg⁠²⁺, Al⁠³⁺, even though their parent atoms come from the same row of the table.

The periodic trends collapse into two practical predictions for any pair of elements:

Bond type prediction (ionic vs covalent). Bond polarity is a continuum, not a set of bright lines: the polarity of a bond increases as the absolute value of the electronegativity difference between the two atoms increases. A bond between identical atoms (ΔEN = 0) is nonpolar covalent; as ΔEN grows the bond becomes polar covalent with the more electronegative atom carrying a partial negative charge (δ−) and the less electronegative atom carrying δ+; at very large ΔEN the electron pair is captured almost entirely by the more electronegative atom and the bond is described as ionic. Common rule-of-thumb cutoffs (often ΔEN < ~0.4 nonpolar, ~0.4–1.7 polar covalent, > ~1.7 ionic) are useful pedagogical anchors but are not absolute thresholds — H–F at ΔEN ≈ 1.8 is typically presented as polar covalent rather than ionic, and many real bonds (AlCl₃, BeCl₂) sit on the borderline and show partial ionic character in some environments and partial covalent character in others.

The qualitative shortcut covers most introductory cases: metal + nonmetal usually means ionic (most chlorides of group-1 and group-2 metals); nonmetal + nonmetal usually means covalent (CO₂, NH₃, CH₄); metal + metal means metallic bonding (treated as a separate category). Use ΔEN to sharpen the prediction at borderlines or to compare polarities between two covalent bonds (e.g., O–H is more polar than C–H because EN(O) > EN(C)).

Property prediction from valence-electron similarity. Elements in any one group (or column) have the same number of valence electrons. Same valence-electron count means similar chemistry: alkali metals (group 1) all form +1 cations and react vigorously with water; halogens (group 17) all form −1 anions; noble gases (group 18) are essentially inert. To identify an unknown element from its chemistry, find the group whose characteristic behavior matches and then place by period. To predict an element’s likely behavior, name its group and apply the family pattern.

Diagonal relationships. A second similarity pattern pairs elements that are not in the same group but lie one period down and one group to the right of each other — the most commonly cited examples are Li and Mg, Be and Al, and B and Si. The two members of a diagonal pair tend to share similar charge-to-radius ratios and similar polarizing power, so their compounds and reactivity patterns track each other more closely than their column neighbors do. Treat diagonal relationships as a secondary similarity rule: same-group similarity is always your first move; diagonal similarity is a second-tier prediction to reach for when a problem points you at one of these specific pairs.

Boundary note: this LO predicts whether a bond will be ionic or covalent from periodic position. The mechanism of ionic bonding (the actual electron transfer that converts neutral atoms into the cation and anion of an ionic compound) is taught and assessed under T3 LO 3.7 — see T3: Atomic Structure and Isotopes.

When you have to rank a mixed set of elements or predict which way a trend will tilt:

1. Locate every species on the periodic table (period, group, block).
2. Apply the trend arrow: down-and-left is bigger / lower IE / lower EN; up-and-right is smaller / higher IE / higher EN.
3. For ions, account for added or removed electrons first, then compare Z (more protons = smaller, when electron count matches).
4. Check for known exceptions: group 13 vs group 2, group 16 vs group 15 in IE; the Cr/Cu electron configurations from Topic 12.
5. Sanity-check with a neighbor — if your prediction puts F below Na in electronegativity, something went wrong.

High-frequency errors: forgetting that ions can shift size by an entire shell (Na vs Na⁠⁺); applying the trend across an exception without flagging it; treating ΔEN cutoffs as exact thresholds rather than guides.