Periodic Trends Study Pack
Kibin's free study pack on Periodic Trends includes a 6-section study guide, 25 quiz questions, 30 flashcards, and 5 open-ended Explain review questions. Sign up free to track your progress toward mastery, plus upload your own notes and recordings to create personalized study packs organized by course.
Last updated May 27, 2026
Periodic Trends Study Guide
Trace how atomic radius, ionization energy, electronegativity, and electron affinity shift across periods and down groups — and why. Master the competing forces of effective nuclear charge and electron shielding that drive every trend, including key exceptions at Groups 2 and 15.
Key Takeaways
- •Atomic radius generally decreases across a period (left to right) due to increasing nuclear charge pulling electrons closer, and increases down a group as additional electron shells are added.
- •Ionization energy is the energy required to remove an electron from a gaseous atom; it increases across a period and decreases down a group, roughly mirroring the atomic radius trend.
- •Electron affinity measures the energy change when a gaseous atom gains an electron; it generally becomes more negative (more energy released) across a period, with notable exceptions at Group 2 and Group 15 due to filled and half-filled subshells.
- •Electronegativity, the tendency of a bonded atom to attract shared electrons, follows the same directional trend as ionization energy — highest at the top-right of the periodic table (fluorine is the most electronegative element) and lowest at the bottom-left.
- •Metallic character increases down a group and decreases across a period, correlating with how easily an atom loses valence electrons.
- •All periodic trends are ultimately explained by two competing factors: effective nuclear charge (Zeff) and electron shielding by inner-shell electrons.
Why Periodic Trends Exist: Effective Nuclear Charge and Shielding
Every periodic trend in atomic properties traces back to the balance between the pull of the nucleus and the buffering effect of electrons already present in the atom — two quantities chemists call effective nuclear charge and shielding.
Effective Nuclear Charge (Zeff)
- •Zeff is the net positive charge experienced by a valence electron after accounting for the repulsion from inner electrons; it is approximated as Zeff = Z − S, where Z is the atomic number and S is the shielding constant.
- •As you move left to right across a period, Z increases by one with each element but the added electrons enter the same principal shell, providing very little additional shielding — so Zeff rises steadily across a period.
- •Because Zeff increases across a period, valence electrons are held progressively more tightly, which drives most of the left-to-right trends.
Electron Shielding and the Effect of Moving Down a Group
- •Core electrons (those in shells below the valence shell) shield valence electrons from the full nuclear charge because their negative charge partially cancels the nucleus's pull.
- •Moving down a group adds a new principal energy level with each period, inserting a new set of core electrons between the nucleus and the outermost electrons — this increases shielding substantially.
- •Even though Z also increases down a group, the shielding effect of the new core electrons roughly offsets the higher nuclear charge, so valence electrons are held less tightly and sit farther from the nucleus.
Atomic Radius: Size Patterns Across and Down the Table
Atomic radius quantifies how large an atom is, typically measured as half the distance between two identical bonded atoms (covalent radius) or as half the distance between adjacent atoms in a metallic crystal (metallic radius) — the exact definition matters less than understanding what drives the size differences.
Trend Across a Period (Left to Right)
- •Atomic radius decreases from left to right across a period because Zeff increases while all valence electrons occupy the same principal shell.
- •The stronger nuclear attraction contracts the electron cloud, pulling electrons closer to the nucleus without adding a new shell.
- •Example: sodium (Na, atomic number 11) has a larger atomic radius (~186 pm) than chlorine (Cl, atomic number 17, ~99 pm), even though both are in Period 3.
Trend Down a Group
- •Atomic radius increases going down a group because each successive element adds a new principal energy level, placing valence electrons farther from the nucleus.
- •The added shielding from new core electrons also reduces the effective pull on the outermost electrons.
- •Example: lithium (Li) has an atomic radius of ~152 pm, while cesium (Cs), directly below it in Group 1, has an atomic radius of ~262 pm.
Isoelectronic Species and Radius
- •When atoms or ions have the same number of electrons (isoelectronic series), radius decreases as nuclear charge increases, because more protons pull the same electron cloud inward.
- •Example: O²⁻, F⁻, Ne, Na⁺, and Mg²⁺ all have 10 electrons; radius decreases from O²⁻ to Mg²⁺ as Z increases from 8 to 12.
Ionization Energy: Removing Electrons from Gaseous Atoms
Ionization energy is the minimum energy required to remove one electron from a neutral gaseous atom, producing a gaseous cation; it is always endothermic and is reported in kJ/mol.
First Ionization Energy (IE₁) Trends
- •IE₁ generally increases across a period because rising Zeff makes valence electrons harder to remove.
- •IE₁ decreases down a group because valence electrons are farther from the nucleus and more shielded, making them easier to remove.
Notable Exceptions to the General Across-Period Trend
- •IE₁ drops from Group 2 to Group 13 (e.g., Mg to Al) because the Group 13 electron is removed from a p subshell, which is higher in energy and less penetrating than the s subshell — so it requires less energy to remove.
- •IE₁ drops from Group 15 to Group 16 (e.g., N to O) because Group 16 must place a fourth electron into an already half-filled p subshell, creating electron-electron repulsion that makes the paired electron easier to remove.
Successive Ionization Energies
- •Each successive ionization energy (IE₂, IE₃, etc.) is larger than the previous one because removing electrons from an increasingly positive ion becomes progressively harder.
- •A dramatic jump in successive ionization energies occurs when the next electron must be removed from a full inner shell (core electrons), signaling that the element has used up all of its valence electrons — this pattern reveals an element's group number.
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Created by Kibin to help students review key concepts, prepare for exams, and study more effectively. This Study Pack was checked for accuracy and curriculum alignment using authoritative educational sources. See sources below.
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What is the approximate atomic radius of sodium (Na) compared to chlorine (Cl), both in Period 3?
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Effective Nuclear Charge and Electron Shielding
Explain effective nuclear charge (Zeff) and electron shielding in your own words. How do these two factors work together, and why are they considered the root cause of most periodic trends?
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