AP Chemistry — Periodic Trends

40 focused cards on atomic and ionic size, ionization energies, electronegativity, electron affinity, shielding and effective nuclear charge, metallic character, and key exceptions. Includes comparison practice and a sourced reference chart on an answer side.

What is effective nuclear charge, Zeff?

The net positive nuclear attraction experienced by an electron after accounting for shielding by other electrons; a useful simplified model is Zeff ≈ Z − S.

What is electron shielding?

Other electrons partially reduce the attraction between a given electron and the nucleus. Core electrons shield valence electrons especially effectively.

Why do many main-group elements in the same group have similar chemistry?

They have similar valence-electron configurations and often the same number of valence electrons.

How does atomic radius generally change across a period from left to right, and why?

It decreases. Nuclear charge rises while added electrons occupy the same principal shell; shielding does not fully offset the stronger attraction.

How does atomic radius generally change down a group, and why?

It increases because valence electrons occupy higher principal energy levels and lie farther from the nucleus, with more core shielding.

Rank Na, Mg, and Al by increasing atomic radius and explain.

Al < Mg < Na. Across period 3, increasing effective nuclear attraction draws the valence shell inward.

Rank Li, Na, and K by increasing atomic radius.

Li < Na < K. Additional occupied shells increase size down group 1.

What is first ionization energy? Write the general process.

Energy needed to remove an electron from a gaseous neutral atom: X(g) → X⁺(g) + e⁻. It is positive because electron removal requires energy.

How does first ionization energy generally change across a period, and why?

It increases as effective nuclear attraction increases and atomic size generally decreases, making an electron harder to remove.

How does first ionization energy generally change down a group, and why?

It decreases. The outer electron is farther from the nucleus and more shielded, so it is generally easier to remove.

Which has the higher first ionization energy, Na or Cl?

Cl. Its valence electrons experience stronger effective nuclear attraction within the same principal shell.

Which has the higher first ionization energy, Li or K?

Li. Its outer electron is closer to the nucleus and less shielded by filled inner shells.

Why is B’s first ionization energy lower than Be’s despite the general across-period trend?

B loses a higher-energy 2p electron, which is easier to remove than Be’s 2s electron. Subshell energy and penetration cause this exception.

Why is O’s first ionization energy lower than N’s?

O has a paired electron in one 2p orbital. Pairing repulsion makes removal easier than from N’s half-filled 2p arrangement.

What analogous first-ionization exceptions occur in period 3?

Al has lower first IE than Mg because Al loses a 3p electron rather than a 3s electron. S has lower first IE than P because paired 3p electrons experience extra repulsion.

Why do successive ionization energies for the same atom increase?

Each removal leaves a more positive ion whose remaining electrons are held more strongly. Removal from an inner shell produces a particularly large increase.

What does a large jump between IE₂ and IE₃ suggest for a main-group atom?

Two relatively accessible valence electrons were removed before a core electron would be removed. This commonly suggests a group 2 atom.

What is electronegativity?

An atom’s tendency to attract shared electrons in a chemical bond. It is distinct from the energy for removing an electron from an isolated atom.

How does electronegativity generally change across a period and down a group?

It generally increases left to right and decreases down a group. Stronger nuclear attraction and smaller size favor attraction for shared electrons.

Which element has the highest electronegativity on the Pauling scale?

Fluorine. Common classroom Pauling tables leave many noble-gas values undefined; do not treat missing values as zero.

Rank C, N, O, and F by increasing electronegativity.

C < N < O < F.

What is electron affinity, using the energy-change convention?

The energy change when a gaseous atom gains an electron: X(g) + e⁻ → X⁻(g). Exothermic attachment has a negative energy change under this convention.

What is the broad electron-affinity pattern, and why is caution needed?

Electron attachment tends to become more favorable toward the halogens across a period, but electron affinity is irregular. Subshell filling, pairing repulsion, and sign conventions prevent a simple universal arrow rule.

Why is electron attachment to Cl more exothermic than attachment to F?

F’s compact 2p shell creates more electron–electron repulsion for the added electron. Cl’s more spacious 3p shell accommodates the extra electron with less crowding.

How does a cation’s radius compare with its neutral atom, and why?

It is smaller. Electron loss reduces repulsion and may remove the entire outer shell; the remaining electrons are attracted more strongly per electron.

How does an anion’s radius compare with its neutral atom, and why?

It is larger. Added electrons increase electron–electron repulsion while nuclear charge stays the same.

What does isoelectronic mean?

Having the same number of electrons. For example, O²⁻, F⁻, Ne, Na⁺, and Mg²⁺ each have 10 electrons.

For an isoelectronic series, what determines the radius trend?

More protons generally means a smaller radius because the same number of electrons feels greater nuclear attraction.

Rank O²⁻, F⁻, Na⁺, and Mg²⁺ from smallest to largest radius.

Mg²⁺ < Na⁺ < F⁻ < O²⁻. All have 10 electrons; greater nuclear charge produces smaller size.

Which is larger, Fe²⁺ or Fe³⁺, and why?

Fe²⁺. For ions of the same element, the higher positive charge generally corresponds to a smaller radius.

How does metallic character generally change across a period and down a group?

It decreases left to right and increases down a group, broadly following how easily atoms lose electrons.

Where are metals, nonmetals, and metalloids generally located?

Metals occupy the left and center; nonmetals occupy the upper right plus hydrogen; metalloids lie near the staircase boundary.

How do typical monatomic charges relate to main-group position?

Group 1 commonly +1; group 2 +2; group 13 often +3; group 15 often −3; group 16 −2; group 17 −1. These are common patterns, not every oxidation state or every compound.

How does alkali-metal reactivity with water generally change down group 1, and why?

It generally increases as the valence electron becomes easier to remove. Actual reaction behavior also depends on physical and experimental conditions.

How does halogen oxidizing ability generally change down group 17?

It generally decreases from F₂ to Cl₂ to Br₂ to I₂. Reaction energetics involve more than electron affinity alone, including bond and solvation energies.

Why should an explanation of periodic trends mention more than just “more protons”?

Changes in shielding, principal shell, distance, subshell energy, and electron pairing can reinforce or offset nuclear charge. State the factor that matters for the comparison.

What directions summarize increasing atomic radius, metallic character, first IE, and electronegativity?

Atomic radius and metallic character generally increase down and left. First IE and electronegativity generally increase up and right, with exceptions. Reference chart: Mirek2, Periodic trends.svg, Wikimedia Commons, CC0; source diagram captured on a white background. Source: https://commons.wikimedia.org/wiki/File:Periodic_trends.svg The chart’s electron-affinity arrow is only a broad tendency toward more favorable attachment; electron affinity has important exceptions and can use different sign conventions.

Why are atomic radius, ionization energy, and electronegativity related but not interchangeable?

All reflect electron–nucleus interactions, but radius measures size, IE measures energy to remove an electron, and electronegativity describes attraction for shared electrons in a bond.

Which has a more strongly attracted valence electron, Na or Na⁺, and why?

Na⁺ holds its remaining electrons much more strongly: removing Na’s 3s electron leaves a filled n=2 shell, and the ion is positively charged. The second ionization requires much more energy than the first.

A main-group atom shows a huge jump after the third electron is removed. What does that suggest?

It likely has three valence electrons, commonly a group 13 atom. The fourth removal would enter the core; interpret with the actual data and configuration.