Chemistry: Atoms First 2e · Representative Metals, Metalloids, and Nonmetals
Periodicity
On this page 9 sections
In 30 seconds
Periodicity The recurring pattern of element properties with increasing atomic number. Full entry → is the observation that the properties of the elements repeat in a regular, predictable pattern as atomic number increases — the reason the periodic table is arranged in rows and columns at all. Atomic radius The size of an atom, roughly half the distance between bonded nuclei. Full entry →, Ionization energy The energy needed to remove an electron from a gaseous atom. Full entry →, Electron affinity The energy change when a gaseous atom gains an electron. Full entry →, and Electronegativity A relative measure of an atom's pull on electrons in a bond. Full entry → all rise and fall across each period and down each group in systematic ways. Underneath these trends sits a single physical cause: the balance between the positive charge of the nucleus and the shielding provided by inner electrons, summarized by the effective nuclear charge (Zeff) experienced by the outermost electrons.
Because the representative elements in Chapter 18 (alkali metals, alkaline earth metals, aluminum, the metalloids, and the nonmetals) span the entire table, periodicity is the organizing framework for everything that follows: why sodium reacts violently with water while chlorine reacts eagerly with sodium, why silicon sits between metal and nonmetal behavior, and why the chemistry of lithium resembles that of magnesium. Master the trends and you can predict reactivity, bonding type, and compound properties for elements you have never studied in detail.
Why this matters
Periodicity is the most powerful prediction tool in descriptive chemistry. A healthcare professional meets its consequences daily: sodium and potassium ions regulate nerve signaling, calcium and magnesium build bone and activate enzymes, and iodine deficiency causes goiter — all behaviors traceable to each element's position in the table. Industrial chemists use trends to choose elements for alloys, semiconductors, and battery materials; geochemists use them to predict which elements concentrate together in ores. On exams, periodic trends are a favorite question type because they test understanding rather than memorization: given two elements, rank their radii, ionization energies, or electronegativities, and justify the ranking with Zeff and shielding arguments.
The college version
Core Concepts
Effective nuclear charge drives the trends
The electrons that determine an element's chemistry are the valence electrons, those in the outermost shell. They are pulled toward the nucleus by its full positive charge but pushed away by inner (core) electrons, which shield the valence electrons from part of the nuclear charge. The net pull is the effective nuclear charge:
Zeff = Z - S
where Z is the atomic number and S is the shielding constant, roughly equal to the number of core electrons. Moving across a period, each added proton adds a unit of nuclear charge while the new electron joins the same shell and shields poorly, so Zeff increases steadily. Moving down a group, each new shell adds many core electrons that shield the valence electrons almost completely, so Zeff increases only slightly — while the principal quantum number n grows, pushing the valence shell farther out.
Atomic radius
Atomic radius is the distance from the nucleus to the outermost electron cloud, usually reported as half the distance between bonded nuclei. Across a period, rising Zeff pulls the same-shell electrons inward, so radius decreases (e.g., Na 186 pm → Cl 99 pm). Down a group, each new shell adds a layer, so radius increases (e.g., F 71 pm → Cl 99 pm → Br 114 pm → I 133 pm), despite the small Zeff increase. Cations are much smaller than their neutral atoms (the atom loses an entire shell or shrinks under the same nuclear charge with fewer electrons), and anions are larger.
Ionization energy
Ionization energy (IE) is the energy required to remove one electron from a gaseous atom: M(g) → M+(g) + e-. First ionization energy increases across a period (electrons are harder to remove as Zeff rises) and decreases down a group (outer electrons are farther away and more shielded). Two well-known exceptions break the smooth climb across a period: the IE1 of boron is lower than beryllium's, because the new 2p electron is higher in energy and more easily removed than a 2s electron; and the IE1 of oxygen is lower than nitrogen's, because the fourth 2p electron must pair with an electron already in an orbital, and the electron–electron repulsion makes it easier to remove. Removing a second electron always costs much more energy than the first (IE2 > IE1), and removing one from a filled shell costs dramatically more — which explains why magnesium forms Mg2+ but sodium only Na+.
Electron affinity and electronegativity
Electron affinity (EA) is the energy change when a gaseous atom gains an electron: X(g) + e- → X-(g). It is generally exothermic (negative) for nonmetals and becomes more negative across a period toward the halogens, which eagerly accept an electron to complete a shell. Electronegativity is a different concept: a relative measure of an atom's ability to attract electrons in a chemical bond. It increases across a period and decreases down a group, with fluorine (3.98 on the Pauling scale) the most electronegative element. The electronegativity difference between bonded atoms predicts bond type: large differences give ionic bonds, small differences give polar covalent bonds, and identical atoms give pure covalent bonds.
Metals, metalloids, and nonmetals
The trends partition the table into three classes. Metals (left and center) have low ionization energies and electronegativities, lose electrons readily, and conduct electricity; they form cations and basic oxides. Nonmetals (upper right) have high ionization energies and electronegativities, gain electrons readily, and are poor conductors; they form anions and acidic oxides. Metalloids (B, Si, Ge, As, Sb, Te — the staircase region) have intermediate properties: silicon's resistivity falls sharply when doped, which makes it the foundation of semiconductor electronics. Across each period the character shifts smoothly from metallic to nonmetallic.
Diagonal relationships
Some pairs of elements in adjacent periods and groups resemble each other more than their group neighbors: Li and Mg, Be and Al, B and Si. The diagonal relationship arises because moving down a group (radius up, IE down) is partly canceled by moving right across a period (radius down, IE up). Lithium and magnesium, for example, both form nitrides when burned in air and react only slowly with water — unlike sodium, which reacts violently. These similarities matter for predicting the chemistry of the lighter elements.
How It Works / Step-by-Step Process
- Locate the elements in the periodic table: identify their period (row) and group (column).
- Compare periods first: the element farther right has the higher Zeff, so it is smaller, harder to ionize, and more electronegative.
- Compare groups second: the element farther down has the larger valence shell, so it is larger, easier to ionize, and less electronegative.
- When the two factors conflict (one element is down and right of another), reason carefully: the dominant factor depends on how far each element is displaced.
- Check for the standard exceptions (Be→B and N→O in IE1) and apply the result to predict bonding or reactivity.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Ionization energy | Electron affinity | IE is the energy to remove an electron; EA is the energy change when an atom gains one. High IE and favorable EA together drive ionic bond formation. |
| Electronegativity | Electron affinity | Electronegativity is an atom's pull on electrons in a bond (a relative scale); EA is the energy of adding a free electron to a gaseous atom. |
| IE trend exceptions | Trend failure | The Be→B and N→O dips are real but have specific subshell explanations (new p subshell; pairing repulsion); the overall trend still holds. |
| Cation size | Anion size | Cations are smaller than their atoms (lost electrons, same nucleus); anions are larger (added electrons, same nucleus). |
| Moving down a group | Moving across a period | Down: new shell → bigger radius, lower IE; across: higher Zeff → smaller radius, higher IE. |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Think of the nucleus as a magnet and the outer electrons as balls on strings. Across a row, the magnet gets stronger and pulls the balls in tighter, so atoms get smaller and hold their electrons harder. Down a column, each atom adds a new, farther-out string, so the balls are looser and easier to pull away — even though the magnet is a bit stronger.
Worked example
Worked example 1: ranking first ionization energies
Rank Na, Mg, and Al by first ionization energy and explain the pattern.
Write the electron configurations first: Na is [Ne]3s1, Mg is [Ne]3s2, and Al is [Ne]3s2 3p1. All three are in period 3, where Zeff rises left to right, so IE1 should generally increase. The measured values follow: Na 496 kJ mol-1, Mg 738 kJ mol-1, Al 578 kJ mol-1. The surprise is that Al is lower than Mg despite being farther right: the new 3p electron sits in a higher-energy subshell and is shielded from the nucleus by the 3s electrons, so it is easier to remove than a paired 3s electron would be. The correct ranking is therefore IE1(Na) < IE1(Al) < IE1(Mg), and the exception is explained by subshell energy and shielding — not by a failure of the trend.
Worked example 2: predicting ion sizes
Which is larger: the potassium ion K+ or the calcium ion Ca2+? Both ions are isoelectronic with argon (18 electrons), so the difference must come from the nucleus. K has Z = 19 and Ca has Z = 20. Write the Zeff comparison: with identical electron configurations, the ion with more protons pulls the same 18 electrons harder:
Zeff(Ca2+) > Zeff(K+)
Therefore Ca2+ is smaller (about 100 pm) than K+ (about 138 pm). This is why potassium and calcium ions behave so differently in biology despite sitting next to each other in the table: Ca2+'s smaller radius and higher charge give it much stronger interactions with phosphate and proteins, letting it serve as a signaling ion that K+ cannot.
Worked example 3: predicting bond type from electronegativity
Predict whether the bonds in NaCl, HCl, and Cl2 are ionic, polar covalent, or pure covalent, using Pauling electronegativities: Na 0.93, H 2.20, Cl 3.16.
Compute each difference before classifying:
Δχ(NaCl) = 3.16 - 0.93 = 2.23
Δχ(HCl) = 3.16 - 2.20 = 0.96
Δχ(Cl2) = 3.16 - 3.16 = 0
A difference above about 1.7 indicates ionic character, values between roughly 0.4 and 1.7 indicate polar covalent bonds, and a difference of zero indicates a pure covalent bond. So NaCl is largely ionic (the electronegativity gap reflects sodium's low IE and chlorine's high EA), HCl is polar covalent with the shared pair pulled toward chlorine, and Cl2 is purely covalent. This single calculation lets you predict solubility behavior, melting points, and conductivity for an entire family of compounds.
Key takeaways
- Zeff increases across a period and barely changes down a group; it is the single cause behind radius, IE, EA, and electronegativity trends.
- Atomic radius decreases across a period, increases down a group; cations are smaller than their atoms, anions larger.
- Ionization energy increases across a period (with Be→B and N→O exceptions) and decreases down a group; IE2 is always much larger than IE1.
- Electronegativity increases across a period and decreases down a group; F is the most electronegative element.
- Metals sit left/low-IE, nonmetals right/high-IE, metalloids on the staircase between them.
- Diagonal relationships (Li–Mg, Be–Al, B–Si) let you predict the chemistry of light elements.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
Why does atomic radius decrease across a period but increase down a group?
Show answer
Across a period, added protons raise Zeff while electrons join the same shell, pulling the cloud inward; down a group, each new shell adds a layer that outweighs the small Zeff increase, pushing electrons farther out.
Explain the two exceptions to the smooth increase in first ionization energy across period 2.
Show answer
Be→B: boron's 3p electron sits in a higher-energy subshell and is easier to remove than beryllium's paired 2s electron. N→O: oxygen's fourth 2p electron must pair with an electron in an already-occupied orbital, and the repulsion makes it easier to remove.
Which is larger, S2- or Cl-, and why?
Show answer
S2-, because both ions are isoelectronic (18 electrons), but sulfur has only 16 protons versus chlorine's 17, so Cl- pulls the electrons harder and is smaller.
Arrange F, Cl, Br, I in order of increasing electronegativity.
Show answer
I < Br < Cl < F; electronegativity decreases down a group as the valence shell grows and shielding increases.
Why do lithium and magnesium resemble each other more than lithium and sodium do?
Show answer
The diagonal relationship: moving down from Li to Na adds a shell and lowers IE, but moving right from Na to Mg raises Zeff — for Li and Mg the two effects nearly cancel, giving similar size, ionization energy, and chemistry.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Periodicity
- The recurring pattern of element properties with increasing atomic number.
- Effective nuclear charge Zeff
- The net positive charge an outer electron feels after core-electron shielding.
- Atomic radius
- The size of an atom, roughly half the distance between bonded nuclei.
- Ionization energy
- The energy needed to remove an electron from a gaseous atom.
- Electron affinity
- The energy change when a gaseous atom gains an electron.
- Electronegativity
- A relative measure of an atom's pull on electrons in a bond.
- Metalloid
- An element with properties between metals and nonmetals.
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
Educational content only. It is not medical, legal or professional advice. Found an error? Tell us.