06 Electron Configurations and Periodic Trends Online Quiz Questions

A neutral oxygen atom has eight electrons. Give its electron configuration and describe one valid distribution of its four 2p electrons among the three 2p orbitals. Explain how the distribution follows the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.

A neutral potassium atom has 19 electrons. Use the stated subshell filling order to write its full electron configuration, then show how noble-gas shorthand represents the same configuration.

Explain how the periodic table’s blocks relate to electron configurations, and how similarities among outer-electron configurations in a main-group column help account for recurring chemical properties.

Why should a simple subshell filling-order prediction not be treated as sufficient to establish the ground-state electron configurations of chromium and copper? Explain the relevant energy consideration without assuming a specific alternative configuration.

As you move left to right across a period, explain why effective nuclear charge generally increases even though electrons are also being added to the atoms.

Use periodic position and electron structure to explain both why sodium is generally larger than chlorine and why potassium is larger than lithium.

Interpret X(g) → X⁺(g) + e⁻ in words and explain what makes the energy associated with this process the first ionization energy.

Predict how first ionization energy generally changes down a group, and explain how the change in energy levels and shielding contributes to that trend.

Boron lies to the right of beryllium, yet its first ionization energy is lower. Explain why this is an exception to the general across-period trend by comparing the subshells of their outermost electrons.

Define electron affinity using the energy-change convention, explain what a more negative value means, and state why the sign convention must be checked when comparing values from different references.

Under the energy-change convention, chlorine’s electron affinity is more negative than fluorine’s. Explain why this comparison can differ from a simple expectation based on their group positions.

A student predicts that electron affinity will become steadily more negative across every period. Evaluate this prediction and explain at least two electron-configuration circumstances that can cause departures from the simple trend.

Oxygen has four electrons in its 2p subshell. Describe how those electrons occupy the three 2p orbitals, and explain how Hund’s rule and the Pauli exclusion principle account for the arrangement.

A neutral chlorine atom is stated to have 17 electrons. A student proposes $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^6$ as its electron configuration. Identify the problem with the proposal and give the correct configuration using the supplied filling order. Show how the electron count supports your correction.

Explain how the periodic table’s blocks and the shared position of elements in a main-group column convey different but related information about electron configurations and chemical behavior.

Across a period, both proton number and electron number increase. Explain why the effective nuclear charge felt by outer electrons nevertheless tends to rise.

An element is replaced by the element directly below it in the same group. Explain the two structural changes that generally make the lower element’s outer electrons less strongly attracted to its nucleus.

Using the location of each element’s outer electron, predict which atom has the larger atomic radius: potassium or lithium. Explain why the energy-level difference supports your prediction.

Define first ionization energy with the corresponding process for a gaseous atom. Then use its general trends across periods and down groups to compare an element near the upper right of the periodic table with one near the lower left.

Beryllium and boron are neighbors in a period, yet boron has a lower first ionization energy. Explain why this comparison departs from the general across-period trend.

Define electron affinity using the process it describes. If a table reports electron affinity as an energy change, explain what a more negative value means and why the table’s convention matters when comparing it with another reference.

With electron affinity expressed as an energy change, chlorine’s value is more negative than fluorine’s. Explain why this comparison is possible even though fluorine is higher in the group, referring to where the added electron enters in each atom.

A student assumes electron affinity must follow a smooth periodic trend because some other properties show general directional patterns. Explain why that assumption can fail, identifying the kinds of electron placements that can disrupt a simple prediction.

Oxygen has the electron configuration $1s^2\,2s^2\,2p^4$. Describe how its four $2p$ electrons occupy the three equal-energy $2p$ orbitals. Explain how Hund’s rule and the Pauli exclusion principle each constrain that arrangement.

Compare the atomic sizes of sodium and chlorine, then potassium and lithium. For each pair, identify which atom is generally larger and explain the comparison using the relevant periodic-table direction and its effect on the electron cloud.

Across a period, first ionization energy generally increases, yet boron’s first ionization energy is lower than beryllium’s. Explain which atom’s outermost electron is easier to remove and why the subshells involved account for this exception.

Electron affinity can be reported as the energy change when an atom gains an electron. Using that convention, compare chlorine’s and fluorine’s electron affinities. Explain the subshell-related reason for the comparison and clarify what a more negative value signifies. What convention issue should you check when comparing values from another reference?

A student assumes that the commonly taught orbital filling order always gives the exact observed ground-state configuration of every element. Use chromium and copper to evaluate this assumption. Explain why exceptions can occur, and state what can and cannot be concluded from the information given here.

A learner claims that an element’s block and its main-group column communicate the same information about its electrons. Explain what the block indicates, what elements in the same main-group column tend to share, and how that shared feature helps account for recurring chemical properties.

Two laboratories report different atomic radii for the same element. One used a covalent-radius measure, while the other used a different defined measure. Explain why the results are not necessarily contradictory and what information is needed to compare them meaningfully.