Ace AP Chemistry Topic 1.7 with this periodic trends practice quiz. Test your knowledge of atomic radius, ionization energy, and Coulomb’s law to prepare for test day.
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The Pre-Quiz Review
Before taking the quiz, make sure you can do more than identify the direction of a periodic trend. AP Chemistry questions often ask you to explain why a trend occurs using ideas such as effective nuclear charge, shielding, electron configuration, and electrostatic attraction. Review the relationships below and practice applying them to unfamiliar elements and ions.
1. Periodic Trends: The Big Picture
Periodic trends describe predictable changes in atomic properties across the periodic table. The most important trends for AP Chemistry are atomic radius, ionization energy, electronegativity, and electron affinity.
Across a Period: Left → Right
- Atomic radius generally decreases.
- First ionization energy generally increases.
- Electronegativity generally increases.
- Electron affinity generally becomes more favorable (more negative), with important exceptions.
As you move across a period, the number of protons increases while added electrons enter the same principal energy level. Because shielding does not increase dramatically within the same shell, the nucleus has a greater attraction for the electrons. This increases the effective nuclear charge (Zeff) experienced by valence electrons.
Down a Group: Top → Bottom
- Atomic radius increases.
- First ionization energy decreases.
- Electronegativity decreases.
- Electron affinity generally becomes less favorable, although electron affinity does not follow as simple a pattern as the other trends.
Moving down a group adds principal energy levels. Valence electrons are farther from the nucleus and experience greater shielding from inner electrons. The attraction between the nucleus and valence electrons therefore decreases.
Key AP Chemistry idea: Do not simply say that a trend occurs because there are “more electrons” or “more protons.” Explain the trend in terms of nuclear charge, shielding, distance from the nucleus, and electrostatic attraction.
2. Effective Nuclear Charge and Shielding
Effective Nuclear Charge (Zeff)
Effective nuclear charge is the net positive attraction experienced by an electron from the nucleus after accounting for shielding by other electrons.
Across a period, Zeff generally increases because nuclear charge increases while shielding remains relatively similar. The stronger attraction pulls electrons closer to the nucleus, contributing to a smaller atomic radius.
Down a group, additional inner energy levels provide more shielding. Although the nucleus has more protons, the valence electrons are farther away and more strongly shielded, so the outer electrons are held less tightly.
What to Mention in an AP Explanation
When explaining a periodic trend, look for this chain of reasoning:
More protons → greater nuclear attraction → larger Zeff → stronger attraction to electrons
or:
More energy levels → greater distance and shielding → weaker attraction to valence electrons
3. Atomic Radius
Atomic radius is a measure of atomic size. There is no single sharp boundary around an atom, so atomic radius is generally determined from measurable distances between bonded atoms.
Trend
- Across a period: atomic radius decreases.
- Down a group: atomic radius increases.
Across a period, electrons are added to the same principal energy level while the nucleus gains protons. Increasing Zeff pulls the electron cloud closer to the nucleus.
Down a group, electrons occupy higher principal energy levels. The valence electrons are farther from the nucleus and experience more shielding, so the atom becomes larger.
Atomic Radius vs. Ionic Radius
Be careful not to confuse atomic radius with ionic radius.
- Cations are smaller than their neutral atoms. Losing electrons decreases electron-electron repulsion and may remove an entire occupied energy level.
- Anions are larger than their neutral atoms. Adding electrons increases electron-electron repulsion.
For isoelectronic species (species with the same number of electrons), the species with more protons is generally smaller because its electrons experience greater nuclear attraction.
For example, O2−, F−, Na+, and Mg2+ all have 10 electrons. Mg2+ is the smallest because it has the greatest nuclear charge.
4. Ionization Energy
Ionization energy is the energy required to remove an electron from a gaseous atom or ion.
The first ionization energy can be represented conceptually as:
X(g) → X+(g) + e−
Trend
- Across a period: first ionization energy generally increases.
- Down a group: first ionization energy generally decreases.
A smaller atomic radius and greater Zeff make an electron more strongly attracted to the nucleus, so more energy is required to remove it.
Important AP Exceptions
The general trend has two famous irregularities that are commonly tested.
- Be has a higher first ionization energy than B. B begins removing an electron from a higher-energy 2p orbital, whereas Be loses a 2s electron.
- N has a higher first ionization energy than O. N has a half-filled 2p3 arrangement. In O, one 2p orbital contains a pair of electrons, creating additional electron-electron repulsion and making removal slightly easier.
The same types of comparisons can appear in other periods when subshell energies and electron-electron repulsions are relevant.
Successive Ionization Energies
Successive ionization energies are the energies required to remove additional electrons:
IE1 < IE2 < IE3 < …
There is often a large jump after all valence electrons have been removed. The large increase indicates that the next electron must be removed from an inner shell, where it experiences much stronger attraction to the nucleus.
AP Chemistry may use this information to determine the likely number of valence electrons in an unknown element.
5. Electronegativity
Electronegativity describes an atom’s ability to attract bonding electrons toward itself in a chemical bond.
Trend
- Across a period: electronegativity generally increases.
- Down a group: electronegativity generally decreases.
Fluorine has the highest electronegativity among the elements. Noble gases are commonly omitted from basic electronegativity trends because they generally do not form bonds under ordinary conditions.
Electronegativity increases toward the upper-right region of the periodic table. This occurs because atoms in that region have relatively small radii and strong attraction for bonding electrons.
Electronegativity and Bond Polarity
The difference in electronegativity between two bonded atoms helps determine bond polarity.
- Similar electronegativities → electrons are shared more equally.
- Different electronegativities → electrons are pulled more strongly toward the more electronegative atom.
The more electronegative atom receives a partial negative charge (δ−), while the other atom receives a partial positive charge (δ+).
Important: Bond polarity and molecular polarity are not always the same. Molecular geometry determines whether individual bond dipoles cancel.
6. Electron Affinity
Electron affinity refers to the energy change associated with adding an electron to a gaseous neutral atom.
X(g) + e− → X−(g)
For many main-group atoms, adding an electron releases energy, so the electron affinity is negative when expressed as an energy change for the process.
General Trend
Electron affinity generally becomes more favorable across a period and less favorable down a group, but it has more irregularities than atomic radius, ionization energy, or electronegativity.
Atoms near the right side of the periodic table, especially halogens, have a strong tendency to gain an electron because doing so moves them closer to a filled valence shell.
Important Exceptions and Ideas
- Noble gases have unfavorable electron affinities because an added electron would have to enter a new principal energy level.
- Group 2 elements have relatively less favorable electron affinities because the added electron must enter a higher-energy p subshell.
- Group 15 elements have less favorable electron affinities than the simple trend might suggest because their p3 configuration is relatively stable.
Do not treat electron affinity as identical to electronegativity. Electron affinity concerns the energy change when an isolated gaseous atom gains an electron, whereas electronegativity describes attraction for bonding electrons in a compound.
7. Coulomb’s Law
Coulomb’s Law describes the electrostatic interaction between charged particles. In chemistry, it provides a foundation for understanding why electrons are attracted to nuclei and why charged particles interact.
The relationship is:
F = k(q1q2)/r2
where F is the electrostatic force, q1 and q2 are the charges, r is the distance between the charges, and k is a proportionality constant.
What the Relationship Means
- Greater charge magnitude → stronger electrostatic force.
- Greater distance → weaker electrostatic force.
- The distance effect follows an inverse-square relationship: doubling r reduces the force to one-fourth.
- Opposite charges attract; like charges repel.
Coulomb’s Law helps explain several periodic trends. A nucleus with greater positive charge can exert a stronger attraction on electrons. Conversely, increasing the distance between the nucleus and an electron weakens their attraction.
8. High-Yield Connections to Remember
Many AP Chemistry questions combine several concepts rather than testing each trend independently.
- Across a period: Zeff increases → atomic radius decreases → ionization energy increases → electronegativity increases.
- Down a group: shielding and energy levels increase → atomic radius increases → ionization energy decreases → electronegativity decreases.
- Stronger nuclear attraction generally means electrons are held more tightly.
- Smaller radius generally means a valence electron is closer to the nucleus.
- More shielding reduces the attraction between the nucleus and an outer electron.
- Cations are smaller than their neutral atoms; anions are larger.
- For isoelectronic species, more protons generally means a smaller radius.
- A large jump in successive ionization energies can reveal the number of valence electrons.
9. Before the Quiz: What You Should Be Able to Do
Before starting the quiz, make sure you can confidently:
- Predict the direction of atomic radius, ionization energy, and electronegativity across periods and down groups.
- Explain periodic trends using Zeff, shielding, electron distance, and Coulombic attraction.
- Compare the sizes of atoms, cations, anions, and isoelectronic species.
- Identify and explain the major first-ionization-energy exceptions involving Be/B and N/O.
- Use successive ionization energies to identify the approximate number of valence electrons.
- Distinguish ionization energy, electron affinity, and electronegativity.
- Predict which atom in a bond is more electronegative and determine the direction of bond polarity.
- Recognize that molecular polarity depends on both bond polarity and molecular geometry.
- Use Coulomb’s Law qualitatively to predict how charge and distance affect electrostatic attraction.
- Explain a trend instead of merely stating its direction.
Final AP Chemistry reminder: When asked “why,” build your answer from the physical causes: nuclear charge → shielding → effective nuclear charge → distance from the nucleus → electrostatic attraction. Connecting these ideas is more important than memorizing isolated arrows on the periodic table.