1.7 Periodic Trends

TLDR

Periodic trends are patterns in atomic properties like atomic radius, ionization energy, electron affinity, and electronegativity that repeat across the periodic table. You can explain every trend using three ideas: effective nuclear charge, the shell model, and Coulomb's law. On the AP Chemistry exam, knowing the direction of a trend matters less than being able to explain why it happens.

AP Periodic Table Trends

On the AP Chemistry periodic table, trends are not just arrows to memorize. Atomic radius, ionization energy, electronegativity, and electron affinity all come from the same reasoning: how strongly the nucleus attracts electrons.

Across a period, effective nuclear charge increases because the number of protons increases while shielding stays about the same. That pulls valence electrons closer. Down a group, added shells place valence electrons farther from the nucleus and increase shielding, so the attraction weakens.

Why This Matters for the AP Chemistry Exam

Periodic trends connect directly to one of the most tested skills in AP Chemistry: predicting and explaining chemical properties using models and theories. You will not get full credit for just saying "atomic radius decreases across a period." You need to tie that pattern to electrostatic attraction, effective nuclear charge, and the number of occupied shells.

This topic builds on electron configuration and photoelectron spectroscopy, and it feeds directly into bonding, ionic compounds, and reactivity later in the course. Both multiple-choice questions and free-response questions in this unit reward students who can link a particle-level cause (more protons pulling on electrons) to a measurable property (smaller radius, higher ionization energy). Get comfortable explaining the "why," and you will handle these questions whether they show up as a graph, a data table, or a written explanation.

Key Takeaways

• Every periodic trend can be explained with effective nuclear charge, the shell model, and Coulomb's law.
• Effective nuclear charge increases across a period because protons increase while shielding stays about the same, pulling electrons in tighter.
• Going down a group adds occupied electron shells, so electrons sit farther from the nucleus and feel a weaker pull.
• Atomic radius, ionization energy, electron affinity magnitude, and electronegativity all move together because they share the same underlying cause.
• A large jump between successive ionization energies signals that you have started removing core electrons, which reveals the number of valence electrons.
• Watch for the small exceptions (group 15 vs 16 ionization energy, fluorine vs chlorine electron affinity) because they show up on the exam.

Foundational Concepts for Periodic Trends

Before memorizing arrows, get the three core ideas down. Every trend traces back to these.

Coulomb's Law and Attraction

Coulomb's law describes the electrostatic attraction between charged particles:

$F_{\text{coulombic}} \propto \frac{q_1 q_2}{r^2}$

The force between the positive nucleus and the negative electrons grows with larger charges and shrinks as distance increases. Opposite charges attract, so a stronger nuclear charge or a shorter distance means a tighter pull on electrons. Almost every trend on this page is a direct consequence of this relationship.

Effective Nuclear Charge

The nucleus holds protons and neutrons. Neutrons are neutral, so protons set the actual nuclear charge (Z). Each electron feels attraction to the nucleus and repulsion from other electrons.

Inner (core) electrons shield outer electrons from the full pull of the nucleus. The charge an outer electron actually experiences is the effective nuclear charge. Qualitatively, effective nuclear charge equals the actual nuclear charge minus the charge shielded by other electrons. You do not need to calculate this on the AP exam, but the concept explains nearly every trend.

• Across a period, protons increase but shielding stays roughly the same, so effective nuclear charge increases.
• Down a group, added inner shells increase shielding, so the outer electrons feel a weaker net pull.

Want a refresher on Coulomb's law and shielding? See the Atomic Structure and Electron Configuration guide.

How the Periodic Table Is Organized

The table is arranged so elements with similar properties line up, which is what makes the trends predictable. It has 18 columns (groups) and 7 rows (periods).

Periods (rows): Moving left to right, atomic number increases by one each step, so each element has one more proton. Elements in the same period have the same number of occupied electron shells. For example, sodium (11 protons) and argon (18 protons) are both in period 3 and both have three occupied shells, but argon's higher proton count gives it a stronger pull on its electrons.

Groups (columns): Moving down a group, the number of occupied shells increases. Elements in the same group share the same number of valence electrons, so they bond in similar ways and have similar chemical properties. Group 18, the noble gases, all have a full outer shell, which makes them stable and generally unreactive. Neon and xenon both have eight valence electrons in their outermost shell, but neon has two occupied shells while xenon has five.

Periodic Trends to Know for AP Chemistry

Apply the three core ideas to each trend below. When you get stuck, return to effective nuclear charge, the shell model, and Coulomb's law.

Atomic Radius

Atomic radius is the distance between an atom's nucleus and its valence electrons.

Across a period (decreases): Moving right, the proton count and effective nuclear charge increase while the number of shells stays the same. The stronger pull draws the valence electrons closer, shrinking the radius. Lithium and fluorine both have two shells, but fluorine's larger nuclear charge makes it smaller.

Down a group (increases): Each step down adds an occupied shell, placing valence electrons farther from the nucleus. In group 1, lithium has two occupied shells while cesium has six, so cesium is much larger.

Ionic Radius

Ionic radius is the distance between an ion's nucleus and its valence electrons.

Cations are smaller than their atoms: When metals lose electrons to form positive ions, shielding and electron-electron repulsion decrease, so the remaining electrons are pulled in tighter. If an atom loses its entire outer shell, the size drop is large.

Anions are larger than their atoms: When nonmetals gain electrons to form negative ions, added electron-electron repulsion pushes the electrons apart, increasing the size.

For an isoelectronic series (species with the same number of electrons), the one with more protons holds those electrons more tightly and is smaller.

Electronegativity

Electronegativity measures how strongly an atom's nucleus attracts shared electrons in a bond.

Across a period (increases): More protons mean a greater nuclear charge, so the nucleus pulls bonding electrons more strongly.

Down a group (decreases): Atoms get larger, so the nucleus sits farther from the bonding electrons and the attraction weakens.

Tip: Fluorine is the most electronegative element (Pauling value 4.0). Use its position to compare other elements.

Ionization Energy

Ionization energy is the energy needed to remove an electron from an atom. Removing the most loosely held electron is the first ionization energy, the next is the second, and so on.

Across a period (increases): Radius decreases and effective nuclear charge increases, so valence electrons are held tighter and need more energy to remove.

Down a group (decreases): Added shells place valence electrons farther out, where they are held loosely and easier to remove.

Things to note:

• Each successive ionization energy is larger than the one before, because removing electrons from a more positive ion (and eventually from inner shells) takes more energy.
• There are exceptions. The first ionization energy of group 15 is higher than group 16. In group 16 elements like sulfur, the electron removed comes from an already paired orbital, and the extra electron-electron repulsion makes it slightly easier to remove than expected.
• Beryllium's first ionization energy is higher than boron's, and magnesium's is higher than aluminum's. The electron removed from boron is a p electron that sits at higher energy and farther out than the s electron in beryllium, so it comes off more easily.

Using Ionization Energies to Find Valence Electrons

A common question gives you successive ionization energies and asks for the number of valence electrons. Look for the large jump.

The big jump happens between I2 and I3, which means the third electron came from a core shell much closer to the nucleus. So this element has 2 valence electrons.

Electron Affinity

Electron affinity is the energy change when an atom in the gaseous state gains an electron.

Across a period: generally more negative (more energy released).

Down a group: generally more positive (less energy released).

A more negative value means more energy is released when the electron is added. Electron affinity is usually negative because atoms tend to release energy when gaining an electron. You can reason about this trend using the same logic as electronegativity.

You might expect fluorine to have the largest magnitude, but chlorine actually does. Fluorine is so small that its electrons are crowded close together, and the added repulsion offsets some of the energy that would otherwise be released.

Number of Unpaired Electrons

The number of unpaired electrons also follows periodic patterns based on how subshells fill. This is useful when you need to predict properties of single atoms or atoms in compounds, alongside size, ionization energy, and electron affinity.

How to Use This on the AP Chemistry Exam

Free Response

When a question says "explain" or "justify," name the cause, not just the direction. A strong answer connects the property to effective nuclear charge, distance from the nucleus (shell model), and the resulting Coulombic attraction. For example: "Chlorine has a higher first ionization energy than sodium because chlorine has more protons and a larger effective nuclear charge, pulling its valence electrons in more tightly, so more energy is required to remove one."

Problem Solving

For successive ionization energy data, find the largest jump. The number of electrons removed before that jump equals the number of valence electrons. The jump marks the point where you start pulling from a core shell that is much closer to the nucleus.

MCQ

Many multiple-choice questions ask you to rank elements by a property or pick the larger or smaller species. Compare proton count and number of occupied shells first. For ions and isoelectronic series, count electrons, then let proton count break the tie (more protons means a smaller species).

Common Trap

Do not stop at the trend direction. Questions often test the exceptions and the reasoning behind them, so be ready to explain group 15 vs 16 ionization energy and the fluorine vs chlorine electron affinity case.

Common Misconceptions

• Atomic number alone controls trends. More protons matter, but you have to account for shielding and the number of occupied shells. Effective nuclear charge, not raw proton count, drives the trends.
• Atomic radius increases across a period because there are more electrons. It actually decreases, because the rising effective nuclear charge pulls the same number of shells in tighter.
• Cations and anions are the same size as their neutral atoms. Cations are smaller (less repulsion, fewer or pulled-in electrons) and anions are larger (more electron-electron repulsion).
• More negative electron affinity is "less." A more negative value means more energy released, which indicates a stronger attraction for the added electron.
• Fluorine has the largest electron affinity magnitude. Chlorine does, because fluorine's small size crowds its electrons and adds repulsion.
• Trends never have exceptions. Group 15 vs 16 ionization energy and the beryllium vs boron case are predictable exceptions you should be able to explain with subshell filling.

zation energy, electronegativity, electron affinity, and patterns in successive ionization energies.

Why does atomic radius decrease across a period?

Atomic radius decreases across a period because effective nuclear charge increases while the number of occupied shells stays the same. The stronger attraction pulls valence electrons closer to the nucleus.

Why does ionization energy increase across a period?

Ionization energy increases across a period because the stronger effective nuclear charge holds valence electrons more tightly, so more energy is needed to remove an electron.

Why does atomic radius increase down a group?

Atomic radius increases down a group because each step adds another occupied electron shell. The valence electrons are farther from the nucleus and more shielded by inner electrons.

How does Coulomb's law explain periodic trends?

Coulomb's law explains that attraction is stronger when charges are larger and distance is smaller. In periodic trends, stronger attraction usually means smaller radius, higher ionization energy, and higher electronegativity.

How do successive ionization energies show valence electrons?

Successive ionization energies show valence electrons through the largest jump. The number of electrons removed before the jump equals the number of valence electrons.

Related AP Chemistry Guides

• Unit 1 Overview: Atomic Structure and Properties
• 1.1 Moles and Molar Mass
• 1.2 Mass Spectra of Elements
• 1.3 Elemental Composition of Pure Substances
• 1.4 Composition of Mixtures
• 1.6 Photoelectron Spectroscopy