Chemistry - Atomic Structure

Define Isotope

Isotopes are ATOMS of the same element with the same proton number but different neutron number, hence different nucleon numbers

Define the term relative isotopic mass

Mass of an atom of isotope is relative to 1/12 the mass of an atom of carbon-12 isotope

What does isotopes of the same element tell you about their chemical and physical properties

They have the same chemical properties (as they have the number number of electrons) but different physical properties (as they have different number of neutrons)

State and explain whether ¹H will deflect in an electric field

¹H has one proton. With atoms being completely stripped of their electrons, ¹H would deflect towards the negatively charged plate as ¹H is positively charged.

Describe two similarity and differences between a 2s and 2p orbital

Similarity

• Both orbitals can hold up a MAXIMUM of 2 electrons

• Both orbitals are in the SAME quantum shell

Difference

• 2s orbital is spherical while 2p orbital is a dumb-bell shape

• 2p orbital is at a higher energy level than the 2s orbital

Why are the electronic configurations for Cr and Cu anomalous?

Electronic configurations with half-filled (Cr) or fully filled (Cu) 3d subshells are unusually STABLE due to the SYMMETRICAL DISTRIBUTION OF CHARGE around the nucleus

Define first ionisation energy

The first ionisation energy is the energy NEEDED to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly charged gaseous cations

What are the factors affecting ionisation energy?

Ionisation energy is affected by the number of electronic shells, nuclear charge by the protons, and the shielding effect by other electrons (repulsive force)

Explain the trend of first ionisation energy down a GROUP

First ionisation energies GENERALLY decreases down a group. Nuclear charge INCREASES. The number of electron shells INCREASES, which results in an INCREASE in shielding effect. Electrostatic attraction between the VALENCE electrons and the nucleus becomes WEAKER. LESSER energy is required to remove the valence electron.

Explain the trend of the first ionisation energy across a PERIOD

First ionisation energies GENERALLY increases across a period. Nuclear charge INCREASES. Electrons are added to the same outermost electron shell and hence the number of electron shells REMAINS THE SAME, and shielding effect remains APPROXIMATELY constant. Electrostatic attraction between the valence electrons and the nucleus becomes STRONGER. More energy is required to remove a valence electron.

Why is the first ionisation energy of Al (577 kJ mol⁻¹) lower than that of Mg (736 kJ mol⁻¹)?

Case 1 anomaly : ns vs np between group 2 and 13 (write down the electronic configurations for the elements mentioned above)

Mg: 1s²2s²3p⁶3s² Al:1s²2s²2p⁶3s²3p¹

Nuclear charge of Al is higher than Mg. However, the 3p electron to be removed is at a HIGHER ENRGY LEVEL and LESS STRONGLY attracted to the nucleus than the 3s electron to be removed from Mg. LESSER energy is required to remove the 3p electron in Al than the 3s electron in Mg.

Why is the first ionisation energy of S (1000 kJ mol⁻¹) lower than that of P (1060 kJ mol⁻¹)?

Case 2 anomaly : paired vs unpaired between group 15 and 16 (write down the electronic configurations for the elements mentioned above)

P: 1s²2s²2p⁶3s²3p³ S: 1s²2s²2p⁶3s²3p⁴

Nuclear charge of S is higher than P. However, the 3p electron to be removed from S is a PAIRED electron while that to be removed from P is an UNPAIRED electron. Due to INTER-ELECTRONIC REPULISION between PAIRED electrons in the p orbital, LESSER energy is required to remove the 3p electron in S than in P.

Define second ionisation energy

The second ionisation energy of an element is the energy needed to remove one mole of electrons from one mole of singly charged gaseous cations to form one mole of doubly charged gaseous cations

Explain the trend in successive ionisation energy of an element

number of protons remains the same, leading to constant nuclear charge. Number of electrons decrease, resulting in a DECREASE in shielding effect. REMAINING electrons experience strong electrostatic attraction by the nucleus. More energy is required to remove an electron, leading to an increase in ionisation energy.

Why is the 3rd IE significantly larger than the 2nd IE for Mg

Mg⁺: 1s²2s²3p⁶3s¹ Mg²⁺: 1s²2s²2p⁶

The 3rd electron is removed from the NEXT INNER electron shell, which is NEARER TO and LESS SHIELDED from the nucleus. Thus, the electron experiences an even stronger electrostatic attraction by the nucleus. EVEN MORE energy is required to remove the electron, leading to a larger increase in IE.

The first 8 ionisation energies of an element X are as follows:
940 2100 3100 4100 7100 7900 15000 16000

State and explain the group of the periodic Table to which X is likely to belong

How to deduce?

Step 1: compute the difference in energy between successive I.E

Step 2: Look out for a SHARP increase in I.E.

Step 3: Find the number of electrons before the SHARP increse in I.E. (this indicates the number of electrons in the valence shell → group number)

Explanation: The largest energy difference is between the 6 and 7 I.E.The 7th electron is MORE strongly attracted to the nucleus as it is removed from the next inner electron shell that is CLOSER to the nucleus and hence LESS SHIELDED from the nucleus.

A student believes that by singly assigning the electrons to the orbital in the valence shell imparts greater stability to the silicon atom because there is no inter-electronic repulsion between the paired electrons.

With reference to the difference in energy level of the 3s and 3p subshells, explain why the student’s explanation is incorrect

The energy required to promote an electron from the 3s subshell to the 3p subshell is greater than the energy arising from the inter-electronic repulsion between the paired electrons in the 3s subshell.