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Quizzes > High School Quizzes > Science

Practice Quiz: The Periodic Table

Sharpen your skills in atomic numbers and periodic law

Difficulty: Moderate
Grade: Grade 10
Study OutcomesCheat Sheet
Paper art promoting Periodic Power Play, a chemistry trivia quiz for high school students.

Which of the following best describes the organization of the periodic table?
Elements arranged by color and shape
Elements arranged alphabetically by element name
Elements arranged by increasing atomic number and grouped by similar chemical properties
Elements arranged in a random order
The periodic table is systematically arranged by atomic number, leading to a periodic repetition of chemical properties. This organization helps in understanding trends and properties of elements.
What does the atomic number of an element represent?
The number of neutrons in the nucleus
The average mass of the element
The number of protons in the nucleus
The total number of electrons in a neutral atom
The atomic number uniquely identifies an element by indicating the number of protons in its nucleus. This number is fundamental to determining the element's chemical behavior.
Which element is represented by the symbol 'O'?
Gold
Oxygen
Oganesson
Osmium
The symbol 'O' stands for oxygen, which is an essential non-metal element. Oxygen plays a critical role in processes like respiration and combustion.
What common property do elements in the same group of the periodic table share?
They have identical complete electron configurations
They have the same atomic mass
They have similar chemical properties
They are always in the same physical state at room temperature
Elements in the same group generally have similar chemical properties due to their similar valence electron configurations. This similarity leads to predictable trends in their reactivity and behavior.
Which of the following elements is a non-metal?
Sodium
Oxygen
Iron
Aluminum
Oxygen is a non-metal characterized by its gaseous state at room temperature and its essential role in many chemical reactions. Metals like sodium, iron, and aluminum display distinctly different properties such as conductivity and malleability.
How does atomic radius change as you move from left to right across a period?
It decreases
It first increases, then decreases
It remains constant
It increases
As you move from left to right across a period, the increasing effective nuclear charge pulls electrons closer to the nucleus. This results in a decrease in atomic radius.
How does increasing effective nuclear charge affect an element's ionization energy?
It has no effect on ionization energy
It decreases ionization energy
It decreases the atomic radius
It increases ionization energy
An increased effective nuclear charge means that the nucleus exerts a stronger pull on the electrons, making them harder to remove. Consequently, the ionization energy increases.
Which element has a higher electronegativity based on periodic trends?
Iodine
Chlorine
Fluorine
Bromine
Fluorine is the most electronegative element because of its small size and high effective nuclear charge. These factors enable it to attract electrons in a bond more strongly than the other halogens.
What property of alkali metals primarily accounts for their high reactivity?
Their large atomic radius
Their dense nucleus
Their single valence electron
Their high electronegativity
Alkali metals have one electron in their outermost shell, which makes it easy to lose that electron and form positive ions. This characteristic is the primary reason behind their high reactivity.
In which period do transition metals (d-block elements) first appear?
Period 4
Period 5
Period 3
Period 2
Transition metals begin to appear in Period 4, marking the start of filling the 3d subshell. This distinguishes them from the s- and p-block elements found in earlier periods.
What electron configuration characteristic is common among noble gases?
They have complete valence electron shells
They have one electron in their outermost shell
They have partially filled d subshells
They have no electrons in their valence shell
Noble gases possess full valence electron shells, which makes them exceptionally stable and unreactive. This complete octet is the reason behind their lack of tendency to form chemical bonds.
As you move down a group in the periodic table, what happens to the atomic radius?
It decreases
It remains constant
It first increases then decreases
It increases
Moving down a group, additional electron shells are added, which increases the atomic radius. Despite the increasing nuclear charge, the effect of shielding causes the outer electrons to be farther from the nucleus.
Which factor has the least impact on an element's placement in periodic trends?
Ionization energy
Electron configuration
Effective nuclear charge
Number of neutrons
The number of neutrons primarily affects the mass and isotope characteristics of an element, but not its chemical properties that dictate periodic trends. Effective nuclear charge and electron configuration play a much larger role in determining an element's behavior.
Which periodic trend best describes the change in electronegativity across a period from left to right?
It generally decreases
It remains constant
It generally increases
It fluctuates unpredictably
Across a period, the effective nuclear charge increases and the atomic radius decreases, which enhances the atom's ability to attract bonding electrons. This results in a general increase in electronegativity from left to right.
What does the periodic law state regarding the properties of elements?
Elements are arranged randomly without any specific pattern
Elements only display trends based on their mass
Elements show recurring properties when arranged by increasing atomic number
Elements in the same period have identical properties
The periodic law asserts that when elements are arranged in order of increasing atomic number, their chemical and physical properties recur periodically. This principle underlies the structure and organization of the periodic table.
Given two elements in the same period with similar atomic numbers, which factor most strongly influences a higher ionization energy?
Greater atomic size
Higher effective nuclear charge
More electrons in inner shells
Higher number of neutrons
Within a single period, a higher effective nuclear charge means that the nucleus exerts a stronger pull on the electrons, making them harder to remove. This increased attraction results in a higher ionization energy. Factors like atomic size or the number of neutrons play a less significant role in this context.
Element X is located in group 17 and period 3. Compared to group 17 elements in period 2, which of the following trends is most likely for Element X?
It will have a larger atomic radius and lower electronegativity
It will have the same atomic radius but higher ionization energy
It will have a smaller atomic radius and higher electronegativity
It will have more valence electrons
As you move down a group, the addition of electron shells increases the atomic radius, while increased shielding results in a decrease in electronegativity. Despite these changes, the number of valence electrons remains constant within a group.
If two elements have similar atomic radii but differing ionization energies, what is the most plausible explanation?
Differences in effective nuclear charge despite similar atomic sizes
One is a transition metal and the other is not
Different numbers of neutrons in their nuclei
Their positions in different periods
Even if two elements have similar atomic sizes, a variation in effective nuclear charge can markedly affect the strength of attraction between the nucleus and the electrons. The element with a higher effective nuclear charge will have a higher ionization energy. The number of neutrons or their classification as transition metals does not significantly account for this difference.
How does the shielding effect influence ionization energy as you move down a group in the periodic table?
Increased shielding increases ionization energy due to more electrons
Shielding increases electronegativity instead of ionization energy
Increased shielding decreases ionization energy despite a higher nuclear charge
Shielding has no impact on ionization energy
Down a group, additional electron shells result in increased shielding, which lowers the effective nuclear charge experienced by the outer electrons. This reduced attraction makes it easier to remove an electron, thereby decreasing the ionization energy. The other options do not accurately reflect the role of the shielding effect.
An element in period 4 and group 15 exhibits a high second ionization energy compared to the first. What best explains this phenomenon?
Electron configuration does not influence ionization energies
Removal of a valence electron makes it significantly harder to remove an inner electron
The first ionization removes all electrons from the outer shell
The element loses two electrons simultaneously
A large jump between the first and second ionization energies indicates that after the removal of a loosely held valence electron, the next electron to be removed belongs to an inner, more tightly bound shell. This results in a significantly higher energy requirement for the second ionization. The phenomenon reflects the stability provided by a full or nearly full valence shell.
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Study Outcomes

  1. Understand the structure and organization of the periodic table.
  2. Identify and interpret periodic trends such as electronegativity, atomic radius, and ionization energy.
  3. Analyze the relationship between element groups and their chemical properties.
  4. Apply periodic trends to predict element reactivity and behavior.
  5. Synthesize information from the periodic table to solve chemical problems in practice scenarios.

Periodic Table Quiz: Atomic Number & Law Cheat Sheet

  1. Atomic Radius - Think of an atom like an inflatable balloon: as you move left to right across a period it shrinks because the nucleus tugs electrons in tighter, and as you drop down a group it swells when new electron shells are added. It's like pulling or adding layers to your favorite party hat! Periodic Trends (Wikipedia)
  2. Ionization Energy - This is the "party-crasher fee" an atom charges to kick out one of its electrons. It generally climbs across a period as the nucleus holds on tighter, and eases down a group since outer electrons are chillin' farther away. Periodic Trends Notes (Save My Exams)
  3. Electronegativity - Picture an atom's craving for electrons in a bond - some are total electron magnets. This craving intensifies across a period and relaxes down a group; fluorine wins the "most clingy" award. Periodic Trends (Wikipedia)
  4. Electron Affinity - This tells you how much an atom loves grabbing an extra electron, shown by the energy change when it happens. Generally the attraction gets stronger (more negative) across a period, showing atoms getting greedier for electrons - though there are a few quirky exceptions. Periodic Trends Video (Learning Box)
  5. Metallic Character - Elements that are happy-go-lucky with electrons - easily letting them go - are called metals. This generosity dips across a period but soars as you move down a group and electrons become loose friends. Periodic Trends Video (Learning Box)
  6. Valency - Think of valency as an atom's number of "friend invites" based on valence electrons. It rises from 1 to 4 across a period, then drops back to zero at the noble gas "VIP-only" gates. Periodic Trends (Wikipedia)
  7. Effective Nuclear Charge (Zeff) - This is the net positive pull a valence electron feels after inner electrons throw some shade. Zeff generally cranks up across a period, making electrons feel more tugged-at, and stays pretty steady down a group. Periodic Trends (Wikipedia)
  8. Shielding Effect - Inner electrons block the nucleus's full pull on outer ones, like VIPs hogging the spotlight. Shielding stays roughly the same across a period but ramps up when you add more shells down a group. Periodic Trends (Wikipedia)
  9. Reactivity of Metals and Nonmetals - Metals get wilder and more reactive down a group, shedding electrons like party confetti. Nonmetals play the opposite game, becoming less eager to grab electrons as you move down. Periodic Trends (Wikipedia)
  10. Mnemonic for Periodic Trends - "Electronegativity, Ionization energy, and Electron affinity increase up and to the right; Atomic radius and Metallic character increase down and to the left." This catchy rhyme is your cheat code for mastering the general trends! Periodic Trends Video (Learning Box)
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