Atomic Structure

Question 1

Define what is first ionisation energy. Why is the sign positive?

Answer 1

First ionisation energy is the energy required to remove one mole of electrons from one more of gaseous atoms to form one mole of singly positively charged gaseous ions.

Sign of 1st and 2nd IE is positive, indicating that it’s an endothermic process. Energy is required to overcome the electrostatic forces of attraction between the nucleus of the gaseous atom and the electron

Question 2

What is nuclear charge and screening effect?

Answer 2

Nuclear charge is the electrostatic attraction between the protons in a nucleus and the surrounding electrons
Screening effect refers to the partial decrease in electrostatic attraction between the nucleus and its valence electrons due to repulsive forces from electrons present.

Question 3

Explain why first IE decreases down the group (similar to AR explanation)

Answer 3

Inc in number of electronic shells, inc in shielding effect.
Valence electrons lie increasingly further away from the nuc, and less strongly attracted to the nucleus, despite increase in nuclear charge. Hence, smaller amt of energy required to remove the electron.

Question 4

Explain why first IE increases across period

Answer 4

Compare nuclear charge and SE
Nuc charge increase due to increase in no of protons
SE remains relatively constant bc electrons added to the valence shell
Zeff increases, stronger EFOA between nucleus and valence electrons
Hence, more energy required to remove the valence electron

Question 5

Explain why electronegativity decreases down the group. How does this differ with the explanation for increasing atomic radius down the group?

Answer 5

Increasing atomic radius:
- Down group, number of electronic shells increase, increasing shielding effect.
VE increasingly further away from nucleus, and less strongly attracted to nucleus despite inc in nuclear charge. Hence, AR increase

Decreasing electronegativity:
Add “ Ability of atom to pull shared pair of e towards itself decreases.”

Down group, number of electronic shells increase, leading to an increase in shielding effect. Valence electrons are increasingly further away from the nucleus. Electrostatic forces of attraction between valence electrons and the nucleus are weaker, despite the increase in nuclear charge. Ability of atom to attract shared pair of electrons towards itself decreases.

Question 6

Explain why there is a general increase in SUCCESSIVE Ie values (eg, 1st ie, 2nd ie, third ie)

Answer 6

Nuclear charge remains unchanged as number of protons remain unchanged. As electrons are removed, shielding effect decreases. The increasingly positive ion atttracts remaining valence electrons more strongly, thus increasing amounts of energy required to remove each remaining VE

Question 7

Explain why there is a SIGNIFICANT increase in ie values from 2nd to third IE (for eg, where graph shows large gap between successive ie values)

Answer 7

3rd ie involves the removal of a 3p electron from an inner shell which experiences greater electrostatic attraction with the nucleus compared to the outermost 4s electron for the 2nd ie. More energy required to remove 3p electron than 4s electron, resulting in higher value for 3rd ie

Question 8

Qn: determine group number if the largest increase in IE is from 2nd to 3rd ie (write out explanation)

Answer 8

Given that the largest increase in IE occurs from 2nd to 3rd ie, the third electron is removed from an inner shell while the second electron is removed from a valence shell. Thus, element has 2 VEs, and is in group 2