S1.3 Electron configurations Questionbank

Write the electron configuration of sodium ($Z = 11$) and magnesium ($Z = 12$).

Deduce the number of core and valence electrons in each atom.

State and explain which atom has a higher first ionization energy.

Magnesium can also form a ${Mg^{2+}}$ ion. Explain how its electron configuration changes and why this ion is stable.

State the atomic number of aluminium.

Write the full electron configuration of a neutral aluminium atom.

State the number of valence electrons in an aluminium atom.

Explain why aluminium tends to form a 3+ ion when it reacts.

State the atomic number of sodium.

Write the full electron configuration of a neutral sodium atom.

Identify the electron that is removed when sodium forms an ion.

State what the boxes and arrows represent in an orbital box diagram.

State the full electron configuration for a neutral atom that would have exactly five electrons filling these orbitals.

Draw the orbital box diagram for the atom described in part 2.

Explain why the 4s orbital is filled before the 3d orbitals when electrons are added to an atom.

State what the energy value of 0 kJ mol^-1 represents for the He⁺ ion.

Calculate the wavelength, in nm, of the photon emitted when an electron falls from $n=4$ to $n=2$. (Planck’s constant $h = 6.63 \times 10^{-34}\ \mathrm{J\ s}$, speed of light $c = 3.00 \times 10^8\ \mathrm{m\ s^{-1}}$)

Determine the first ionization energy of the He⁺ ion in $\mathrm{kJ}\ \mathrm{mol}^{-1}$ .

Compare the energy required for the electron transitions from $n=2$ to $n=3$ for the He⁺ ion and the H atom.

Sketch a diagram showing the emission spectrum for He⁺ corresponding to transitions ending at $n=1$ (Lyman series), indicating the trend in line spacing.

Discuss why the emission lines converge at high frequencies, and how this convergence relates to ionization.

State the maximum number of electrons that can be placed in the $2p$ orbitals and explain their spin arrangement according to Hund's rule.

The orbital box diagram corresponds to an element in its ground state. Deduce the identity of this element if all the orbitals shown are fully occupied.

Draw the orbital box diagram for the element identified in Part 2, clearly indicating electron spins.

Compare the first ionization energies of this atom and its neighbor with one fewer proton. Explain your answer in terms of nuclear charge and electron shielding.

The energy levels of the 2s and 2p orbitals are not exactly the same. Explain why the 2p orbital is slightly higher in energy than the 2s orbital.

Define the term electron configuration.

Write the full electron configuration for a phosphorus atom ($Z = 15$).

Phosphorus is in Group 15. Explain how its electron configuration supports its group placement.

Suggest why phosphorus forms a ${P^{3-}}$ ion and describe the resulting electron configuration.

Write the electron configuration of aluminum ($Z = 13$).

Identify the number of unpaired electrons in a ground-state aluminum atom.

Explain the difference between the electron configurations of a neutral aluminum atom and an ${Al^{3+}}$ ion.

Explain the position of aluminum in the periodic table using its electron configuration.

State the maximum number of electrons that can be accommodated in the $2 p$ sublevel.

Outline the significance of the arrows that would be drawn inside the boxes in a full orbital box diagram.

State the full electron configuration for a neutral atom that fills all the orbitals shown in the diagram.

Draw the orbital box diagram for the atom described in part 3, clearly showing the spin of electrons.

State and explain why electrons fill each $2 p$ orbital singly before pairing up.

State the number of electrons in a neutral neon atom.

Write the full electron configuration of neon.

Describe what happens to an electron in an atom when it absorbs energy.