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Quizzes > High School Quizzes > Science

Periodic Table STAAR Practice Quiz

Master element facts with interactive quiz questions

Difficulty: Moderate
Grade: Grade 8
Study OutcomesCheat Sheet
Colorful paper art promoting a Periodic Table Power trivia quiz for high school chemistry students.

Which element is an alkali metal?
Uranium
Lithium
Helium
Carbon
Lithium is an alkali metal found in Group 1 of the periodic table. Helium, Carbon, and Uranium belong to different groups with distinct properties.
What is the general trend observed as you move from left to right across a period?
Atomic weight always decreases
Atomic radius decreases
Atomic radius increases
Number of electron shells increases
As you move across a period, electrons are added to the same shell and the nucleus pulls them closer, resulting in a decrease in atomic radius. The other options do not accurately describe the observed trend.
What is the trend for ionization energy across a period?
Ionization energy decreases
Ionization energy first increases and then decreases
Ionization energy remains constant
Ionization energy increases
Ionization energy increases across a period because increasing effective nuclear charge makes electron removal more difficult. The other options are not consistent with periodic trends.
Which element is known for its inertness due to a full valence electron shell?
Neon
Aluminum
Sodium
Chlorine
Neon is a noble gas with a complete valence electron shell, which makes it chemically inert. In contrast, sodium, aluminum, and chlorine are more reactive elements.
What is the primary basis for the modern arrangement of elements in the periodic table?
Increasing atomic weight
Increasing atomic number
Decreasing atomic radius
Alphabetical order
The periodic table is arranged in order of increasing atomic number, which reflects the unique structure and properties of each element. The other options do not serve as the organizing principle of the periodic table.
Which periodic trend best explains why fluorine is more electronegative than chlorine?
Fluorine is in a higher period than chlorine
Fluorine has more electron shells than chlorine
Chlorine has a filled d-subshell
Fluorine has a smaller atomic radius and higher effective nuclear charge
Fluorine's small atomic size combined with a high effective nuclear charge allows it to attract electrons more effectively, making it highly electronegative. The other options do not correctly describe the factors influencing electronegativity.
What is the general trend for metallic character as you move across a period?
Metallic character remains unchanged
Metallic character increases
Metallic character decreases
Metallic character first increases then decreases
Across a period, elements tend to become less metallic and more nonmetallic due to increasing effective nuclear charge. This results in a decrease in metallic character from left to right.
Which of the following elements is most likely to form a +1 ion?
Aluminum
Sodium
Magnesium
Chlorine
Sodium belongs to Group 1, where elements typically lose one electron to form a +1 ion. Magnesium and aluminum form ions with higher positive charges, while chlorine typically forms a -1 ion.
What trend is observed for electron affinity as you move from left to right across a period?
Electron affinity generally increases
Electron affinity remains constant
Electron affinity generally decreases
Electron affinity decreases then increases
As atoms approach a complete electron configuration across a period, they tend to gain electrons more readily, leading to an increase in electron affinity. The other options do not reflect the typical periodic trend.
Which element typically has the smallest atomic radius in period 2?
Beryllium
Neon
Fluorine
Lithium
In period 2, atomic radius decreases from left to right, making neon, with the highest effective nuclear charge, have the smallest atomic radius among the listed elements. The other elements have relatively larger radii due to their positions in the period.
Why do elements in the same group often exhibit similar chemical properties?
They all have the same atomic mass
They have identical nuclear charges
They are in the same period
They have similar numbers of valence electrons
Elements in the same group have similar valence electron configurations, which largely determine their chemical behavior. The other options are inaccurate because atomic mass, period, and nuclear charge vary among group elements.
What is the effect of increased shielding on the atomic radius of an element?
Increased shielding results in a smaller atomic radius
Increased shielding typically leads to a larger atomic radius
Shielding has no effect on atomic size
Shielding decreases the effective nuclear charge to zero
When shielding increases, the outer electrons experience a reduced effective nuclear charge, which allows them to reside further from the nucleus, thereby increasing the atomic radius. The other statements either contradict this effect or are oversimplified.
Which property increases as you move up a group in the periodic table?
Atomic radius increases
Electron affinity decreases
Metallic character increases
Ionization energy increases
Moving up a group leads to a decrease in atomic size and shielding, which makes it more difficult to remove an electron, thereby increasing ionization energy. The other options either describe opposite trends or are not consistently observed.
How does effective nuclear charge affect the attraction between the nucleus and electrons?
There is no relationship between effective nuclear charge and electron attraction
Effective nuclear charge only affects inner electrons
Greater effective nuclear charge weakens electron attraction
Greater effective nuclear charge strengthens electron attraction
A higher effective nuclear charge means a stronger positive charge experienced by the electrons, pulling them closer to the nucleus. This directly enhances the attraction between the nucleus and the electrons.
Which element is most likely to have the highest ionization energy among the following?
Neon
Potassium
Sodium
Magnesium
Neon, a noble gas, has a complete electron shell and a high effective nuclear charge, which makes it very difficult to remove an electron. The other elements, being more reactive metals, have lower ionization energies.
How does the shielding effect influence the trends in ionization energy down a group?
Shielding effect has no impact on ionization energy
Increased electron shielding down a group increases ionization energy
Increased electron shielding down a group decreases ionization energy
Ionization energy is solely determined by atomic number
As you move down a group, additional electron shells increase the shielding effect, which reduces the effective nuclear charge felt by the outer electrons. This weaker attraction makes it easier to remove an electron, thus lowering the ionization energy.
Which of the following best explains why fluorine's electron affinity is slightly lower than chlorine's?
Fluorine's compact electron cloud increases electron-electron repulsions when gaining an electron
Both fluorine and chlorine have identical electron affinities due to similar valence electron configurations
Fluorine's higher electronegativity always results in the highest electron affinity among halogens
Electron affinity decreases with atomic number in halogens, making fluorine the lowest
Although fluorine is highly electronegative, its small size leads to increased electron-electron repulsions when an extra electron is added. This effect slightly lowers its electron affinity compared to chlorine, which has a larger atomic radius.
How does the concept of effective nuclear charge explain the periodic trend observed in atomic radii across a period?
Effective nuclear charge remains constant across a period, so it does not influence atomic radius
Increasing effective nuclear charge pulls electrons closer, thus reducing atomic radius
Effective nuclear charge only affects inner electrons, not those in the outer shell
Increasing effective nuclear charge pushes electrons away, leading to larger atomic radii
As the effective nuclear charge increases across a period, electrons are drawn closer to the nucleus, which results in a smaller atomic radius. The other options do not accurately describe the impact of effective nuclear charge on atomic size.
Which statement best describes the trends observed for ionization energy and electron affinity in the periodic table?
Both decrease across a period as atomic radii decrease
Both increase across a period due to increasing effective nuclear charge
Ionization energy decreases as electron affinity increases across a period
Ionization energy increases while electron affinity decreases across a period
Both ionization energy and electron affinity tend to increase across a period because the effective nuclear charge increases, making it harder to remove an electron and more favorable to add one. This combined trend reflects the strengthening grip of the nucleus on the electrons.
Why do elements in the same period display a gradual change in properties rather than sudden shifts?
Because the periodic table is arranged by atomic number without considering electron configuration
Because elements in the same period have identical chemical properties
Because electrons fill different orbitals randomly
Because electrons are added to the same energy level, causing a gradual change in effective nuclear charge
Elements in the same period add electrons to the same principal energy level, leading to incremental changes in effective nuclear charge and atomic properties. This results in a gradual transition in chemical and physical properties across the period.
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Study Outcomes

  1. Identify periodic trends such as atomic radius, ionization energy, and electronegativity.
  2. Analyze the relationship between element properties and their positions on the periodic table.
  3. Apply group and period concepts to predict chemical behaviors of elements.
  4. Compare and contrast the properties of metals, nonmetals, and metalloids.

Periodic Table STAAR Quiz Review Cheat Sheet

  1. Atomic Radius - Think of this as the "personal bubble" around an atom. As you zip across a period, the nucleus pulls electrons in tighter, shrinking the radius, while sliding down a group adds new electron shells and puffs up that bubble. Learn more on Wikipedia
  2. Ionization Energy - This is the energy it takes to kick an electron out of its atomic home. Across a period, atoms clutch their electrons more fiercely, raising the energy bar, but down a group, electrons lounge farther from the nucleus and are easier to remove. Learn more on Wikipedia
  3. Electronegativity - Picture this as an atom's magnetic charm for electrons in a bond. It gets stronger left to right as atoms vie for electron attention, and weakens as you go down a group. Fun fact: fluorine is the ultimate electronegativity champion! Learn more on Wikipedia
  4. Electron Affinity - This measures the energy change when an atom eagerly grabs an extra electron. It becomes more negative (more eager) across a period but mellows out going down a group. It's like atoms get more "electron-hungry" as you move rightward. Learn more on Wikipedia
  5. Metallic Character - These are the traits that make metals so handy: conductivity, shine, and bendability. Moving down a group, metallic behavior cranks up, while cruising across a period tames it. In short, heavyweights in the same column are the most metal! Learn more on Wikipedia
  6. Valency - Valency is an element's "friendliness" or how many bonds it can form. Across a period it climbs from 1 up to 4 then tapers down to 0 at the noble gases, and down a group it generally stays true to its value. It's the perfect guide to predicting who teams up with whom! Learn more on Wikipedia
  7. Effective Nuclear Charge (Z_eff) - This is the actual pull an electron feels from the nucleus after accounting for all the other electrons in the way. Z_eff grows stronger across a period, tightening that pull, but remains pretty steady down a group. It's chemistry's secret tug-of-war! Learn more on Wikipedia
  8. Shielding Effect - Inner electrons act like a bodyguard, shielding outer electrons from the nucleus's full force. The more shells you stack going down a group, the stronger the shield, which swells atomic size and lowers ionization energy. Think of it as electrons hiding behind their buddies! Learn more on Wikipedia
  9. Reactivity of Metals and Non-Metals - Metal atoms get more eager to lose electrons (and react) as you slide down a group, while non-metals grow less keen to grab electrons. This explains why cesium explodes with water and chlorine mellows out compared to fluorine. It's chemistry's ultimate popularity contest! Learn more on Wikipedia
  10. Mnemonic for Periodic Trends - Remember: "Across a period: radius decreases, ionization energy and electronegativity increase. Down a group: radius increases, ionization energy and electronegativity decrease." This catchy phrase will keep your periodic trend trivia on-point every time! Learn more on Wikipedia
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