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Quizzes > High School Quizzes > Science

Periodic Trends Practice Quiz

Strengthen your chemistry skills with practice

Difficulty: Moderate
Grade: Grade 10
Study OutcomesCheat Sheet
Paper art representing trivia quiz on Periodic Trends Unleashed for high school chemistry students.

What happens to atomic radius as you move from left to right across a period?
It remains constant
It first increases then decreases
It increases
It decreases
Atomic radius decreases across a period due to an increase in nuclear charge, which pulls electrons closer to the nucleus. This results in a smaller atomic size from left to right.
Which trend is observed for ionization energy as you move from left to right in a period?
Ionization energy remains constant
Ionization energy first increases then decreases
Ionization energy decreases
Ionization energy increases
Ionization energy increases across a period because electrons are held more tightly by the nucleus as a result of increasing nuclear charge. This makes it harder to remove an electron from elements on the right side of the period.
How does electronegativity change as you move from left to right across a period?
It fluctuates randomly
It increases
It stays the same
It decreases
Electronegativity increases from left to right across a period due to the rising effective nuclear charge. This causes atoms to attract electrons more strongly as you move across the period.
What trend is observed in metallic character as you move from left to right across a period?
Metallic character fluctuates
Metallic character decreases
Metallic character increases
Metallic character remains unchanged
Metallic character decreases across a period because elements transition from metals to nonmetals. This change is largely due to increasing electronegativity and a stronger effective nuclear charge.
Why do elements on the right side of the periodic table generally have higher ionization energies?
Their valence electrons are closer to the nucleus
They have increased electron shielding
They have more electron shells
They are larger in size
Elements on the right have higher ionization energies because their valence electrons are held closer to the nucleus by a stronger effective nuclear charge. This makes it more difficult to remove an electron.
How does an increase in effective nuclear charge affect atomic radius across a period?
It increases the atomic radius by repelling electrons
It has no effect on the atomic radius
It decreases the atomic radius by pulling electrons closer
It only affects the nucleus size
An increased effective nuclear charge pulls electrons closer to the nucleus, thereby reducing the atomic radius. This is a fundamental reason why atoms become smaller across a period.
What is the trend in electron affinity as you move from left to right in a period?
Electron affinity becomes positive
Electron affinity remains unchanged
Electron affinity becomes more negative
Electron affinity becomes less negative
As one moves from left to right, electron affinity generally becomes more negative, meaning atoms release more energy upon gaining an electron. This is due to the stronger pull of the nucleus on the additional electron.
Which factor is primarily responsible for the decrease in atomic radius as one moves down a group?
Decrease in electron shielding
Decrease in effective nuclear charge
Increase in electronegativity
Addition of more electron shells
Moving down a group, additional electron shells are added, which increases the distance between the nucleus and the outer electrons. This results in a larger atomic radius despite an increase in nuclear charge.
When comparing isoelectronic species, which statement best describes the impact of nuclear charge on their size?
Higher nuclear charge results in a smaller ionic radius
Nuclear charge has no impact on size
Lower nuclear charge results in a smaller ionic radius
Higher nuclear charge results in a larger ionic radius
For isoelectronic species, an increased nuclear charge pulls the electron cloud closer, thereby decreasing the ionic radius. This concept is crucial when comparing ions with identical electron numbers.
What best explains the decrease in ionization energy as one moves down a group?
Increased electron shielding reduces the pull of the nucleus on outer electrons
Effective nuclear charge increases significantly
Atomic radius decreases
Electrons are closer to the nucleus in lower periods
Down a group, the addition of electron shells leads to increased shielding, which weakens the attraction between the nucleus and the outer electrons. This makes it easier to remove an electron, thus lowering the ionization energy.
How does atomic structure contribute to higher ionization energies across a period?
Greater effective nuclear charge binds electrons more tightly
Larger atomic sizes mean electrons are further away
More electron shielding causes electrons to be easily removed
Increasing number of electron shells leads to higher ionization energy
As you move across a period, the effective nuclear charge increases, which holds the electrons more tightly. This stronger attraction makes it more difficult to remove an electron, resulting in higher ionization energies.
What is the trend in metallic character as you move from left to right across a period?
Metallic character increases
Metallic character decreases
Metallic character first increases then decreases
Metallic character remains constant
Metallic character decreases across a period as elements evolve from metals to nonmetals. This change is influenced by increasing electronegativity and effective nuclear charge.
How does increased electron shielding down a group affect atomic size?
It decreases atomic size by pulling electrons closer
It causes atomic size to fluctuate
It has no significant effect on atomic size
It increases atomic size by reducing nuclear attraction
Increased electron shielding down a group reduces the effective pull of the nucleus on the outer electrons. Consequently, electrons are held less tightly, and the atomic radius increases.
Why do nonmetals typically exhibit higher electronegativities than metals?
Nonmetals have more electron shells
Nonmetals have a stronger effective nuclear charge relative to their atomic size
Nonmetals remove electrons easily
Nonmetals have lower electron affinity
Nonmetals usually have smaller atomic radii and a higher effective nuclear charge relative to their size, enabling them to attract electrons more strongly. This is why they tend to have higher electronegativities compared to metals.
How can electron configuration explain deviations from expected periodic trends in ionization energy or electron affinity?
Only atomic mass affects these deviations
Elements with half-filled or fully-filled subshells exhibit extra stability, affecting these trends
The number of protons is the sole factor influencing these trends
Electron configuration has no impact on periodic trends
Electron configurations can lead to extra stability, especially when subshells are half-filled or fully-filled. This additional stability can cause deviations from the expected trends in ionization energy and electron affinity.
In comparing a period 3 element and a period 4 element from the same group, which statement best describes their trends in atomic radius and ionization energy?
The period 3 element has a larger atomic radius and lower ionization energy
The period 3 element has a smaller atomic radius but lower ionization energy
The period 3 element has a smaller atomic radius and higher ionization energy
Both elements have identical atomic radii and ionization energies
Elements in an earlier period have fewer electron shells, which leads to a smaller atomic radius and higher ionization energy compared to elements in later periods. This is due to reduced electron shielding and a stronger effective nuclear charge in period 3 elements.
What is the primary reason for the lower than expected electron affinity of nitrogen?
The electron configuration leads to a stable half-filled p subshell
Nitrogen has a lower nuclear charge than expected
High electron shielding in nitrogen reduces electron affinity
Nitrogen atoms are too large
Nitrogen has a half-filled p subshell which offers extra stability. This stable configuration means that adding an electron does not release as much energy as expected, resulting in a lower electron affinity.
Which factor best explains the anomaly in the first ionization energy of oxygen compared to its neighboring elements?
Oxygen's electrons are in a lower energy level
Strong electron-electron repulsion in paired p orbitals
Higher effective nuclear charge than expected
Oxygen exhibits increased electron shielding
Oxygen experiences significant repulsion between electrons in its paired p orbitals, which reduces the energy required to remove an electron. This electron-electron repulsion results in an ionization energy that is slightly lower than expected.
How do effective nuclear charge and increased electron shielding collectively influence electronegativity down a group?
They cause electronegativity to increase down a group
They cause electronegativity to remain constant
They lead to a decrease in electronegativity down a group
They have no effect on electronegativity
Although nuclear charge increases down a group, the effect of increased electron shielding and a larger atomic radius diminishes the nucleus's pull on electrons. This results in a decrease in electronegativity for elements further down the group.
Which statement best describes the impact of subshell electron configuration on periodic trends such as ionization energy?
Ionization energy is solely determined by atomic number
Electron configuration anomalies, like half-filled subshell stability, can lead to higher ionization energies than predicted
Subshell electron configuration only affects atomic mass, not ionization energy
Electron configuration has minimal impact on ionization energy trends
Certain electron configurations, especially those involving half-filled or fully-filled subshells, offer extra stability to an atom. This stability makes it more difficult to remove an electron, leading to ionization energies that deviate from the trend predicted by effective nuclear charge alone.
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Study Outcomes

  1. Analyze trends in atomic radii, ionization energy, electron affinity, and electronegativity.
  2. Apply periodic trends to predict element behavior and properties.
  3. Evaluate the relationship between electronic structure and periodic properties.
  4. Explain the variations in chemical reactivity across different periods and groups.
  5. Compare and contrast key periodic trends to enhance understanding of element characteristics.

Periodic Trends Quiz Review Cheat Sheet

  1. Atomic Radius - Think of atoms as having personal space: moving left to right, the nucleus pulls electrons tighter like fans squeezing toward the stage, so the radius shrinks; going down a group adds more electron shells like extra stadium tiers, so the radius grows. Learn more on Wikipedia
    2. en.wikipedia.org/wiki/Periodic_trends
  2. Ionization Energy - This is the energy you need to "steal" an electron, and it gets harder across a period because the nucleus holds on tighter; down a group electrons are farther out and easier to pluck away, so ionization energy drops. Learn more on Wikipedia
    3. en.wikipedia.org/wiki/Periodic_trends
  3. Electronegativity - Imagine atoms playing tug-of-war for electrons: across a period their grip gets stronger and they win more often; down a group their pull weakens because added shells put distance between nucleus and bonding electrons. Learn more on Wikipedia
    4. en.wikipedia.org/wiki/Periodic_trends
  4. Electron Affinity - This measures how happy an atom is to gain an electron: as you move across a period, atoms generally become more eager (more negative affinity), while down a group the extra shells dull their enthusiasm (less negative). Learn more on Wikipedia
    5. en.wikipedia.org/wiki/Periodic_trends
  5. Metallic Character - Metals love sharing electrons around like party favors - down a group they get more "metallic" (shiny, conductive, malleable), but across a period they lose those traits and behave more like shy non‑metals. Learn more on Wikipedia
    6. en.wikipedia.org/wiki/Periodic_trends
  6. Valency - Valency is how many friends (electrons) an atom can bond with: across a period it climbs from 1 to 4 then slides back to 0 at noble gases, while down a group most elements keep the same bonding count. Learn more on Wikipedia
    7. en.wikipedia.org/wiki/Periodic_trends
  7. Effective Nuclear Charge (Zeff) - This is the real "pull" an outer electron feels after inner electrons get in the way; across a period it strengthens as protons accumulate, but down a group added shells keep Zeff nearly steady. Learn more on Wikipedia
    8. en.wikipedia.org/wiki/Periodic_trends
  8. Shielding Effect - Inner electron layers act like bodyguards blocking the nucleus from outer electrons: as you descend a group, more shells mean stronger shielding and less nuclear grip on valence electrons. Learn more on Wikipedia
    9. en.wikipedia.org/wiki/Periodic_trends
  9. Reactivity of Metals and Non‑Metals - Metals get rowdier down a group (easier to lose electrons, think alkali fireworks!), while non‑metals calm down - they hold onto electrons less tightly when more shells separate them from the nucleus. Learn more on Wikipedia
    10. en.wikipedia.org/wiki/Periodic_trends
  10. Mnemonic for Periodic Trends - Remember "BEAR": Basicity increases up & left, Electronegativity/Electron affinity/Ionization energy up & right, Acidity down & right, Radius down & left. It's a fun way to unpack all those trends at once! Check out the full mnemonic guide
    11. prospectivedoctor.com/mcat-mnemonics-periodic-trends
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