Electronic Configurations (DP IB Chemistry)

When electrons return to their original energy levels, they emit energy.

False.

The energy emitted can be in the infrared, visible, or ultraviolet regions of the spectrum.

Electron jumps from n∞ to n1 correspond to the ultraviolet region of the spectrum.

Electron jumps from n∞ to n2 correspond to the visible region of the spectrum.

Electron jumps from n∞ to n₃ correspond to the infrared region of the spectrum.

Convergence is when the spectral lines get closer together towards the higher energy end of the spectrum.

True.

The emission spectrum of hydrogen in the visible region shows lines converge at a higher frequency.

The line emission spectrum of the hydrogen atom provides evidence for the existence of electrons in discrete energy levels which converge at higher energies.

Ionisation energy is the maximum amount of energy an electron can reach, corresponding to the energy required for the electron to escape the atom.

A represents the ionisation of hydrogen in the ground state because the ionisation energy corresponds to a transition from the ground state to n=∞

False.

The higher the principal quantum number, the greater the energy of the electron within that shell.

The following orbital is the p_x orbital.

The following orbital is the p_z orbital.

The following orbital is the p_yorbital.

The following orbital is the s orbital.

True.

The subshells increase in energy as follows: s ＜ p ＜ d ＜ f

The maximum number of electrons that can be held in the n=1 principal energy level is 2.

The maximum number of electrons that can be held in the n=2 principal energy level is 8.

The maximum number of electrons that can be held in the n=3 principal energy level is 18.

The maximum number of electrons that can be held in the n=4 principal energy level is 32.

False.

Each orbital can hold a maximum number of 2 electrons.

The p subshell hold 6 electrons.

The completed electron configuration for sodium is:

1s² 2s² 2p⁶ 3s¹

The completed electron configuration for aluminium is:

1s² 2s² 2p⁶ 3s² 3p¹

The completed electron configuration for calcium is:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

The condensed electronic configuration for gallium is:

[Ar] 3d¹⁰ 4s² 4p¹

Chlorine has the full electronic configuration of 1s² 2s² 2p⁶ 3s² 3p⁵.

The condensed electron configuration for copper is:

[Ar] 3d¹⁰ 4s¹

The condensed electronic configuration for chromium is:

[Ar] 3d⁵ 4s¹

True.

The 4s subshell fills first and empties before the 3d subshell.

False.

[Ar] 3d⁴ 4s² is less energetically favourable than [Ar] 3d⁵ 4s¹

The full electronic configuration of Sr²⁺ is:

1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶

The electronic configuration silicon is:

The electronic configuration using box notion for vanadium is:

ΔE is change in energy.

ΔE = h ν

h is Planck's constant.

ΔE = h ν

v is frequency.

ΔE = h ν

c is the speed of light.

c = ν λ

λ is wavelength.

c = ν λ

The ionisation energy for one atom of hydrogen is:

• ΔE = h ν

• E = 32.883 x 10¹⁴ x 6.63 x 10^-34 = 2.18 x 10^-18 (J)

The equation to calculate the frequency of the convergence limit is:

ν = c ÷ λ

To calculate the energy change per mole (kJ mol^-1):

• IE₁ = 8.22 × 10⁻¹⁹  × 6.02 × 10²³

• IE₁ = 494 844 J mol⁻¹

• IE₁ = 495 kJ mol⁻¹

True.

Down a group ionisation energy decreases as shielding effect increases, therefore, the attraction of outer electrons to the nucleus decreases.

Ionisation energy increase across a period because the nuclear charge increases and the shielding by inner shell electrons remains the same.

Successive ionisation energies increase because removing electrons from a positive ion becomes more difficult due to increasing attractive forces and decreasing shielding.

There is a slight decrease in IE₁ between beryllium and boron as the fifth electron in boron is in the 2p subshell, which is further away from the nucleus than the 2s subshell of beryllium.

True.

There is a slight decrease in IE₁ between nitrogen and oxygen due to spin-pair repulsion in the 2p_x orbital of oxygen.

False.

The increase in successive ionisation energies is not constant and depends on the atom's electronic configuration.

The element is located in Group 13.

There is a large decrease in ionisation energy between periods due to increased distance between the nucleus and outer electrons, and increased shielding by inner electrons in the new shell, which outweigh the increased nuclear charge.

Successive ionisation energy data can be used to:

• Predict or confirm the simple electronic configuration of elements

• Confirm the number of electrons in the outer shell

• Deduce the group an element belongs to in the Periodic Table

The element is found in Group 2 as the largest jump in energy is between Group 2 and Group 13.

• IE₁ = 899 kJ mol^-1,

• IE₂ = 1757 kJ mol^-1,

• IE₃ = 14850 kJ mol^-1,

• IE₄ = 21005 kJ mol^-1