Unlock hundreds more features
Save your Quiz to the Dashboard
View and Export Results
Use AI to Create Quizzes and Analyse Results

OR
Sign inSign in with Facebook
Sign inSign in with Google
Quizzes > Science > Chemistry

Periodic Trends Quiz: Challenge Yourself!

Ready to master periodic table trends? Take the quiz!

Difficulty: Moderate
2-5mins
Learning OutcomesCheat Sheet
Paper art atoms periodic table icons on yellow background for atomic radii ionization energy electron configuration quiz

This periodic trends quiz helps you practice AP Chemistry skills on atomic radius, ionization energy, and electron configuration across groups and periods. Work through mixed questions to build speed and spot gaps before the exam; if you want another option, try the alternate set .

In which direction does atomic radius generally decrease on the periodic table?
From right to left across a period
From left to right across a period
From bottom to top up a group
From top to bottom down a group
As you move from left to right across a period, the effective nuclear charge increases while shielding remains relatively constant, pulling electrons closer and reducing atomic radius. This trend is a fundamental periodic property. Greater nuclear attraction draws valence electrons inward.
Which of the following second period elements has the smallest atomic radius?
Lithium (Li)
Beryllium (Be)
Boron (B)
Carbon (C)
Atomic radius decreases across a period as increasing nuclear charge pulls electrons closer. Among Li, Be, B, and C, carbon has the highest effective nuclear charge with minimal added shielding, resulting in the smallest radius. Shielding by inner electrons remains similar across these elements, so the nucleus more effectively attracts valence electrons.
Which of the following elements has the highest first ionization energy?
Phosphorus (P)
Silicon (Si)
Magnesium (Mg)
Aluminum (Al)
Ionization energy generally increases across a period due to rising effective nuclear charge. Phosphorus, being further to the right in the same period, has a higher first ionization energy than Mg, Al, or Si. The greater nuclear attraction requires more energy to remove an electron.
Which of the following elements is the most electronegative?
Silicon (Si)
Phosphorus (P)
Sulfur (S)
Chlorine (Cl)
Electronegativity increases across a period and decreases down a group. Chlorine is in the same period as sulfur and phosphorus but is further right, making it the most electronegative of the group. Electronegativity reflects an element's tendency to attract bonding electrons.
Which of the following elements has the most exothermic (most negative) first electron affinity?
Iodine (I)
Fluorine (F)
Bromine (Br)
Chlorine (Cl)
Electron affinity generally becomes more exothermic across a period; chlorine has a slightly more exothermic first electron affinity than fluorine due to lower electron - electron repulsion in its larger valence shell. Electron affinity measures the energy change when an atom gains an electron, and a more exothermic value indicates a greater tendency to accept an electron. This trend reflects increasing nuclear attraction for the added electron.
Which element has the electron configuration [Ne] 3s2 3p2?
Silicon (Si)
Phosphorus (P)
Argon (Ar)
Sulfur (S)
[Ne] 3s2 3p2 corresponds to silicon, which has atomic number 14. The configuration shows two electrons in the 3p subshell, matching group 14. This pattern follows the Aufbau principle and the Madelung rule.
Which of the following elements has the highest second ionization energy?
Beryllium (Be)
Carbon (C)
Nitrogen (N)
Boron (B)
Beryllium's second ionization energy is very high because removing a second electron would disrupt the stable 1s2 electron core. First ionization removes a valence 2s electron, but the second would break into the filled 1s2 subshell. This requires substantially more energy than removal from B, C, or N.
Which of the following ions has the smallest ionic radius?
Sodium ion (Na+)
Aluminum ion (Al3+)
Fluoride ion (F?)
Magnesium ion (Mg2+)
In the isoelectronic series, higher positive charge leads to stronger attraction of electrons to the nucleus, shrinking the radius. Al3+ has the greatest nuclear charge for the same electron count and thus the smallest ionic radius. This trend demonstrates the effect of charge on ionic size.
Which statement best describes the trend in first electron affinity across the periodic table?
It becomes more exothermic from top to bottom down a group
It becomes less exothermic from left to right across a period
It becomes more exothermic from left to right across a period
It becomes less exothermic from top to bottom down a group
First electron affinity generally becomes more exothermic across a period due to increased nuclear attraction for additional electrons, though there are some exceptions. Down a group, added electron shells reduce attraction, making affinity less exothermic. Electron affinity trends reflect the nucleus's pull on incoming electrons and can show anomalies at filled or half-filled subshells.
Which of the following is expected to have the largest ionic or atomic radius?
Sulfide ion (S2?)
Argon (Ar)
Potassium ion (K+)
Chloride ion (Cl?)
All species are isoelectronic with 18 electrons, but S2 - carries the highest negative charge, resulting in greater electron - electron repulsion. This repulsion pushes electrons further apart, giving S2 - the largest radius. Isoelectronic trends illustrate how charge affects size in species with equal electron counts.
Which of these chalcogens has the most exothermic first electron affinity?
Selenium (Se)
Sulfur (S)
Oxygen (O)
Tellurium (Te)
Electron affinity generally becomes more exothermic across a period but can vary down a group due to size and repulsion effects. Sulfur, being larger than oxygen, reduces inter-electron repulsion when an extra electron is added. This balance of nuclear pull and lower repulsion makes sulfur's first electron affinity the most exothermic among the chalcogens listed.
Which element has a first ionization energy lower than that of the element immediately before it in the same period due to electron pairing repulsion?
Boron (B)
Nitrogen (N)
Oxygen (O)
Beryllium (Be)
Nitrogen has a half-filled 2p subshell, which is relatively stable and holds electrons more tightly. In oxygen, adding a second electron to a 2p orbital introduces repulsion between paired electrons, increasing their energy and lowering the ionization energy. This extra repulsion makes it easier to remove an electron from oxygen than from nitrogen.
Which element in the third period experiences the greatest effective nuclear charge (Zeff)?
Phosphorus (P)
Sodium (Na)
Argon (Ar)
Magnesium (Mg)
Effective nuclear charge increases across a period as protons are added without full shielding by additional electrons. Argon, at the end of the third period, has the highest nuclear charge with relatively constant shielding, giving it the greatest Zeff. This stronger pull results in smaller atomic size and higher ionization energies.
Among the isoelectronic series O2?, F?, Ne, Na+, and Mg2+, which has the smallest radius?
Oxygen ion (O2?)
Magnesium ion (Mg2+)
Fluoride ion (F?)
Neon (Ne)
In an isoelectronic series, ions with greater positive charge have stronger electrostatic attraction for the same electron count. Mg2+ carries the highest positive charge among these species, pulling electrons closer. As a result, Mg2+ has the smallest radius.
What is the primary factor responsible for the decrease in atomic radius across a period?
Decreased number of electron shells
Increased electron - electron repulsion
Increased effective nuclear charge
Increased electron shielding
Across a period, electrons occupy the same shell but more protons are added to the nucleus, increasing the effective nuclear charge. This stronger nuclear attraction pulls electrons closer, reducing atomic radius. The result is a consistent decrease in size from left to right.
An element X has a first ionization energy of 578 kJ/mol and a second of 1817 kJ/mol. Which element is X?
Silicon (Si)
Aluminum (Al)
Magnesium (Mg)
Phosphorus (P)
The relatively low first ionization energy indicates removal of a valence electron, while the much higher second energy indicates breaking into a noble gas core. Aluminum's known IE1 and IE2 values match these energies. The large increase shows the stability of the neon-like configuration that remains after two ionizations.
What is the term for the phenomenon where the ionic radii of lanthanides decrease steadily from La3+ to Lu3+, affecting the atomic sizes of subsequent elements?
Actinide expansion
Skutterudite effect
D-block contraction
Lanthanide contraction
Lanthanide contraction refers to the steady decrease in ionic radii across the lanthanide series due to poor shielding by 4f electrons. Because 4f electrons shield poorly, added protons draw the remaining electrons closer to the nucleus. This results in smaller-than-expected atomic and ionic radii for subsequent period 6 elements.
Which element has the electron configuration [Ar] 3d10 4s2 4p2?
Gallium (Ga)
Selenium (Se)
Germanium (Ge)
Arsenic (As)
[Ar] 3d10 4s2 4p2 corresponds to an atomic number of 32. This configuration fills the 3d subshell and places two electrons in 4p, characteristic of germanium. It follows the Aufbau principle in the fourth period.
An element has successive ionization energies (kJ/mol): IE1 = 737, IE2 = 1450, IE3 = 7733. Which element is it?
Silicon (Si)
Magnesium (Mg)
Phosphorus (P)
Aluminum (Al)
The large jump between the second and third ionization energies indicates removal of a core electron after the valence electrons are removed. Magnesium's known successive ionization energies match these values. The significant increase reflects the stability of the neon core that remains after two ionizations.
Why is the first ionization energy of beryllium higher than that of boron, deviating from the general trend across period 2?
Beryllium has a larger atomic radius than boron
Beryllium has a filled 2s subshell which is particularly stable
Boron has extra shielding from its 2p electron
Boron has higher electron affinity than beryllium
Beryllium's filled 2s subshell is particularly stable, requiring more energy to remove an electron compared to boron's singly occupied 2p electron. Boron's first ionization removes an electron from a higher-energy p orbital, making it easier. This subshell stability leads to the anomalous trend in first ionization energies.
0
{"name":"In which direction does atomic radius generally decrease on the periodic table?", "url":"https://www.quiz-maker.com/QPREVIEW","txt":"In which direction does atomic radius generally decrease on the periodic table?, Which of the following second period elements has the smallest atomic radius?, Which of the following elements has the highest first ionization energy?","img":"https://www.quiz-maker.com/3012/images/ogquiz.png"}

Study Outcomes

  1. Understand Atomic Radii Trends -

    Explain how atomic size changes across periods and down groups in the periodic table.

  2. Analyze Ionization Energy Patterns -

    Identify the factors that influence the energy required to remove electrons from atoms.

  3. Interpret Electron Configurations -

    Relate electron arrangements to an element's position and chemical properties within the periodic table.

  4. Predict Chemical Reactivity -

    Use periodic table trends to forecast how elements will behave in reactions.

  5. Apply Periodic Trends Quiz Skills -

    Leverage your knowledge of atomic radii, ionization energy, and electron configurations to answer quiz questions accurately.

  6. Evaluate Quiz Performance -

    Assess your results to pinpoint strengths and areas for improvement in AP Chemistry periodic trends.

Cheat Sheet

  1. Atomic Radii Patterns -

    Atomic radii decrease across a period as increasing proton count pulls electrons closer, and increase down a group as electrons occupy higher energy levels. For an atomic radii quiz, remember LAWD ("Largest Atoms go to the Left and Down") to recall these trends. For instance, sodium (181 pm) is larger than magnesium (160 pm) in period 3, a fact often highlighted in AP Chemistry periodic trends resources from university sites like MIT OpenCourseWare.

  2. Ionization Energy Trends -

    First ionization energy generally increases across a period and decreases down a group, reflecting how tightly an atom holds its outermost electron. Use the mnemonic IERUPT ("Ionization Energy Rises Upward and to the Right") when practicing an ionization energy quiz. For example, the energy to remove one electron from magnesium (737 kJ/mol) is higher than from sodium (496 kJ/mol), as detailed on reputable sites such as the Royal Society of Chemistry.

  3. Electron Configuration Stability -

    Atoms achieve extra stability when they have half-filled or fully filled subshells, a concept critical in both AP Chemistry periodic trends and periodic table trends quiz questions. Chromium's configuration [Ar] 4s¹3d❵ illustrates this exception, as explained in peer-reviewed journal articles on electron configuration anomalies. Master these patterns to predict anomalies on your next periodic trends quiz.

  4. Effective Nuclear Charge (Zeff) -

    Effective nuclear charge (Zeff) equals the actual nuclear charge (Z) minus the shielding constant (S), and it explains why ionization energies and radii shift across the table. For oxygen, Zeff ≈ 8 − 2 = 6, demonstrating stronger pull on valence electrons than in nitrogen - details available in University of California resources. Understanding Zeff helps you tackle advanced questions in any AP Chemistry periodic trends review.

  5. Metallic vs. Nonmetallic Character -

    Metallic character increases down a group and decreases across a period, while nonmetallic character shows the opposite pattern; a key theme in any periodic table trends quiz. Alkali metals like cesium react vigorously with water, whereas halogens like fluorine gain electrons to form anions, as highlighted by the American Chemical Society. Recognizing these reactivity patterns will boost your confidence in self-paced periodic trends quizzes.

Make a Quiz (FREE) Copy & Edit This Quiz Create a Quiz with AI

Chemistry Quizzes

Powered by: Quiz Maker